Iron Sulfate: Formulas, Main Uses and Precautions

He Iron sulphate Is a crystalline solid, greenish or yellow-brown in color. In nature, it is in the form of iron (II) sulfate (also known as ferrous sulfate, green caparrosa, green vitriol, among others) and iron (III) sulphate (also called ferric sulphate, Mars vitriol, Among others), each in different degrees of hydration.

It is used for the treatment of water or waste water and as an ingredient of fertilizers. Its main problem is the threat to the environment. Immediate measures should be taken to limit their spread to the environment.

Iron (II) sulphate heptahydrate (formula: FeSO4 7H2O) crystallizes as green monoclinic crystals.

When heated to 60-70 ° C, 3 moles of water are expelled and the iron (II) sulfate tetrahydrate (formula: FeSO 4 4H 2 O) is formed.

Upon heating to about 300 ° C, in the absence of air, a white powder is formed consisting of iron (II) sulfate monohydrate.

When heated to about 260 ° C, and in the presence of air, the monohydrate is oxidized to iron (III) sulfate.

In its anhydrous form, iron (III) sulfate (formula: Fe 2 (SO 4) 3) is a yellowish white solid, which is hydrolyzed by dissolving in water to produce a brownish solution.

• Formulas

	Iron sulphate (II)	Iron sulphate (II)	Iron sulphate (III)
	(anhydrous)	(Heptahydrate)	(anhydrous)
Formula	FeSO4	FeSO4 7H2O	Fe2 (SO4) 3

• CAS : 7720-78-7 Anhydrous iron (II) sulphate
• CAS : 7782-63-0 Iron (II) sulphate heptahydrate
• CAS : 10028-22-5 Iron (III) anhydrous sulphate

2D structure

3D structure

characteristics

Physical and chemical properties

	Iron sulphate (II)	Iron sulphate (II)	Iron sulphate (III)
	anhydrous	Heptahydrate	anhydrous
Appearance	White crystals	Blue-green crystals	Grayish white powder, or crystals
Molecular weight:	151,901 g / mol	278,006 g / mol	399.858 g / mol
Boiling point:	90 ° C	90 ° C
Melting point:	64 ° C		480 ° C
Density:	1898 kg / m3		1898 kg / m3
Solubility in water, g / 100 ml at 20 ° C:	29.5 g / L water		Soluble

Iron (II) sulfate belongs to the group of weak reducing agents. It is a yellow-brown or greenish crystalline solid. The appearance and odor vary depending on the iron salt. The most common form is heptahydrate, blue-green.

Iron sulphate (III) belongs to the group of acid salts. It comes in the form of grayish white powder, or yellow rhombohedral crystals.

Inflammability

• Many weak reducing agents are flammable or combustible. However, they may require extreme conditions (eg high temperatures or pressure) to combust.
• Iron (II) sulfate is non-flammable but, like other weak inorganic reducing agents, it reacts with oxidizing agents to generate heat and products that can be flammable, combustible or reactive.
• None of the acid salts is highly flammable.

Reactivity

• Reactions of weak reducing agents with oxidizing agents can lead to combustion and may be potentially explosive if the mixture is heated or subjected to pressure.
• Oxygen, which is a moderately strong oxidizing agent and is ubiquitous in the atmosphere, can react with the compounds of this class in the presence of disturbances, such as heat, a spark, the action of a catalyst or a mechanical shock.
• Iron (II) sulfate is efflorescent in dry air. In moist air, the surface of the crystals is covered with basic iron (III) brownish sulphate.
• Aqueous solutions of iron (II) sulphate are slightly acidic due to hydrolysis.
• Acid salts react as weak acids to neutralize the bases. These neutralizations generate heat, but less than that generated by the neutralization of inorganic acids, inorganic oxoacids or carboxylic acids.
• Iron (III) sulphate is soluble in water. It is slowly hydrolyzed in aqueous solutions. Form acidic aqueous solutions. It is hygroscopic in the air. It is corrosive to copper, copper alloys, mild steel and galvanized steel.

Toxicity

• Most weak reducing agents are toxic by varying ingestion. They may also cause chemical burns if inhaled or in contact with the skin.
• If ingested, iron (II) sulfate may cause disturbances in the gastrointestinal tract. Ingestion of large amounts by children can cause vomiting, hematemesis, liver damage and peripheral vascular collapse.
• With regard to acid salts, their toxicity is also highly variable. The solutions of these materials are generally corrosive to the skin and irritants to the mucous membranes.
• Inhalation of iron (III) sulphate powder irritates the nose and throat. Ingestion causes irritation of the mouth and stomach. The powder irritates the eyes and can irritate the skin on prolonged contact.

Applications

• Iron (II) sulfate is used for the preparation of other iron compounds.
• It is used in the elaboration of iron inks and pigments, in the process of engraving and lithography, in wood preservatives and as an additive to fodder, among others.
• In these applications, by-products of industrial processes that affect the environment are generated. This has led to other uses being sought for iron (II) sulfate.
• Large amounts of iron (II) sulfate are used to clarify effluent from communities. Sludge formed in clarification tanks can be used as fertilizers.
• The conversion of iron (II) sulfate into gypsum and iron (II) chloride has also been proposed by treatment with calcium chloride.
• As an additive to the cement, iron (II) sulfate can substantially reduce the content of water-soluble chromates.
• Iron (II) sulfate can be used to combat chlorosis, a disease of the vines. It is also used to treat alkaline soil and to destroy moss.
• Iron (III) sulphate is used to prepare aluminas and pigments of iron oxide, and as a coagulant for the treatment of liquid effluents.
• Ferrous ammonium sulfate is used for tanning. Solutions of iron (III) compounds are used to reduce the volume of sludge from effluent treatment plants.

Clinical effects

Iron has historically been one of the leading causes of death from intoxication in children. Exposure has declined in recent years with improved packaging, but still has significant morbidity and mortality.

Iron is required for the normal functioning of essential proteins and enzymes, including hemoglobin, myoglobin and cytochromes, but is a poison to cells and is corrosive to the gastrointestinal mucosa.

It is found as a nutritional supplement in vitamins (generally in the form of iron (II) sulfate or ferrous sulfate). It is used for the treatment and prevention of iron deficiency anemia.

Among the main adverse effects of its therapeutic use are gastrointestinal disorders and constipation.

Symptoms of mild or moderate intoxication include vomiting and diarrhea, within 6 hours after ingestion.

Symptoms of severe intoxication include severe vomiting and diarrhea, lethargy, metabolic acidosis, shock, gastrointestinal bleeding, coma, Convulsions , Hepatotoxicity, and late gastrointestinal stenosis.

Excessive long-term ingestion of iron-containing compounds can lead to increased iron accumulation in the body, particularly in the liver, spleen and lymphatic system, accompanied by fibrosis of the pancreas, diabetes mellitus, and liver cirrhosis. Signs and symptoms may include irritability, nausea or vomiting, and normocytic anemia.

Safety and Risks

Hazard statements of the Globally Harmonized System of Classification and Labeling of Chemicals (GHS).

The Globally Harmonized System of Classification and Labeling of Chemicals (GHS) is an internationally agreed system, created by the United Nations and designed to replace the various classification and labeling standards used in different countries through the use of globally consistent criteria.

The hazard classes (and their corresponding GHS chapter), classification and labeling standards, and recommendations for iron (II) sulfate are as follows (European Chemicals Agency, 2017, United Nations, 2015, PubChem, 2017) :

(United Nations, 2015, p.366). (United Nations, 2015, p.371). (United Nations, 2015, p.382). (United Nations, 2015, p.385).

The hazard classes (and their corresponding GHS chapter), classification and labeling standards, and recommendations for iron (II) sulphate heptahydrate are as follows (European Chemicals Agency, 2017, United Nations, 2015, PubChem, 2017 ):

(United Nations, 2015, p.371). (United Nations, 2015, p.382). (United Nations, 2015, pp. 385).

The hazard classes (and their corresponding GHS chapter), classification and labeling standards, and recommendations for iron (III) sulfate are as follows (European Chemicals Agency, 2017, United Nations, 2015, PubChem, 2017) :

(United Nations, 2015, p.366). (United Nations, 2015, p.371). (United Nations, 2015, p.381). (United Nations, 2015, p.382). (United Nations, 2015, p. 388). (United Nations, 2015, pp. 384). (United Nations, 2015, pp. 385). (United Nations, 2015, p.395).

References

1. Benjah-bmm27, (2007). Sample of iron (II) sulfate heptahydrate [image] Recovered from Wikipedia.org .
2. European Chemicals Agency (ECHA). (2017). Summary of Classification and Labeling.
3. Annex VI of Regulation (EC) No 1272/2008 (CLP Regulation, Diiron tris (sulphate).) Recovered on 16 January 2017, from Echa.europea.eu .
4. European Chemicals Agency (ECHA). (2017). Summary of Classification and Labeling.
5. Annex VI of Regulation (EC) No 1272/2008 (CLP Regulation). Iron (II) sulfate. Recovered on January 16, 2017, from Echa.europea.eu .
6. European Chemicals Agency (ECHA). (2017). Summary of Classification and Labeling.
7. (II) sulfate (1: 1) heptahydrate, sulfuric acid, iron (II) salt (1: 1), heptahydrate
8. Ferrous sulfate heptahydrate. Recovered on January 16, 2017, from Echa.europea.eu .
9. Hazardous Substances Data Bank (HSDB). TOXNET. (2017). Ferrous sulfate. Bethesda, MD, US: National Library of Medicine.
10. Jmol: an open source Java viewer for three-dimensional chemical structures, (2017). Ferrous sulfate.
11. Jmol: an open source Java viewer for three-dimensional chemical structures, (2017). Ferric sulfate.
12. United Nations (2015). Globally Harmonized System of Classification and Labeling of Chemicals (SGA) Sixth Revised Edition. New York, USA: United Nations publication. Recovered from Unece.org .
13. National Center for Biotechnology Information. PubChem Compound Database. (2016). Ferric sulfate - PubChem Structure [image] Recovered from Nhi.gov .
14. National Center for Biotechnology Information. PubChem Compound Database. (2016). Ferrous sulfate - PubChem Structure [image] Recovered from Nhi.gov .
15. National Center for Biotechnology Information. PubChem Compound Database. (2016). Ferrous sulfate - PubChem Structure [image] Recovered from Nhi.gov .
16. National Center for Biotechnology Information. PubChem Compound Database. (2017). Iron (II) sulfate. Bethesda, MD, US: National Library of Medicine. Recovered from Nhi.gov .
17. National Center for Biotechnology Information. PubChem Compound Database. (2017). Iron (II) sulfate heptahydrate. Bethesda, MD, US: National Library of Medicine. Recovered from Nhi.gov .
18. National Center for Biotechnology Information. PubChem Compound Database. (2017). Ferric sulfate. Bethesda, MD, US: National Library of Medicine. Recovered from Nhi.gov .
19. National Oceanic and Atmospheric Administration (NOAA). CAMEO Chemicals. (2017). Chemical Datasheet. Ferrous sulfate. Silver Spring, MD. EU.
20. National Oceanic and Atmospheric Administration (NOAA). CAMEO Chemicals. (2017). Chemical Datasheet. Ferric sulfate. Silver Spring, MD. EU.
21. National Oceanic and Atmospheric Administration (NOAA). CAMEO Chemicals. (2017). Reactive Group Datasheet. Reducing Agents, Weak. Silver Spring, MD. EU.
22. National Oceanic and Atmospheric Administration (NOAA). CAMEO Chemicals. (2017). Reactive Group Datasheet. Salts, Acidic. Silver Spring, MD. EU.
23. Ondřej Mangl, (2007). Iron (III) sulfate [image] Recovered from Wikipedia.org .
24. Siran železnatý, (2007). FeSO4 [image] Recovered from Wikipedia.org .
25. Smokefoot, (2016). Structure of iron (II) sulfate heptahydrate [image] Recovered from Wikipedia.org .
26. Wikipedia. (2017). Iron (II) sulfate. Retrieved January 17, 2017, from Wikipedia.org .
27. Wikipedia. (2017). Iron (II) sulfate. Retrieved January 17, 2017, from Wikipedia.org .
28. Wildermuth, E., Stark, H., Friedrich, G., Ebenhöch, F. L., Kühborth, B., Silver, J., & Rituper, R. (2000). Iron Compounds. In Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH Verlag GmbH & Co. KGaA.