- Structure
- Nomenclature
- Properties
- Physical state
- Molecular weight
- Melting point
- Boiling point
- Density
- Solubility
- pH
- Chemical properties
- Aqueous SO solutions
- Other properties
- Obtaining
- Presence in nature
- Applications
- In the production of sulfuric acid
- In the processed food industry
- As a solvent and reagent
- As a reducing agent
- In various applications
- Effects of OS
- Risks
- Ecotoxicity
- Effects of ingesting it with food
- References
The sulfur dioxide is a gaseous inorganic compound consisting of sulfur (S) and oxygen (O), and its chemical formula SO 2. It is a colorless gas with an irritating and suffocating odor. In addition, it is soluble in water, forming acid solutions. Volcanoes expel it into the atmosphere during eruptions.
It is part of the biological and geochemical cycle of sulfur, but it is produced in large quantities by certain human activities such as oil refining and the burning of fossil fuels (coal or diesel for example).
Sulfur dioxide SO 2 is emitted by volcanoes during eruptions. Brocken Inaglory. Source: Wikimedia Commons.
SO 2 is a reducing agent that allows the paper pulp to remain white after bleaching with other compounds. It also serves to remove traces of chlorine in water that has been treated with this chemical.
It is used to preserve some types of food, to disinfect containers where the fermentation of grape juice is produced to produce wine or barley to make beer.
It is also used as a fungicide in agriculture, to obtain sulfuric acid, as a solvent and as an intermediate in chemical reactions.
The SO 2 present in the atmosphere is harmful to many plants, in the water it affects fish and it is also one of those responsible for "acid rain" which corrodes materials created by humans.
Structure
The sulfur dioxide molecule is symmetrical and forms an angle. The angle is due to the fact that SO 2 has a lone pair of electrons, that is, electrons that do not form a bond with any atom but are free.
Lewis structure of sulfur dioxide where its angular shape and the pair of free electrons are observed. WhittleMario. Source: Wikimedia Commons.
Nomenclature
- Sulfur dioxide
- Sulfur anhydride
- Sulfur oxide.
Properties
Physical state
Colorless gas.
Molecular weight
64.07 g / mol
Melting point
-75.5 ºC
Boiling point
-10.05 ºC
Density
Gas: 2.26 at 0 ° C (relative to air, that is, air density = 1). This means that it is heavier than air.
Liquid: 1.4 to -10 ° C (relative to water, that is, density of water = 1).
Solubility
Soluble in water: 17.7% at 0 ° C; 11.9% at 15 ° C; 8.5% at 25 ° C; 6.4% at 35 ° C.
Soluble in ethanol, diethyl ether, acetone, and chloroform. It is less soluble in non-polar solvents.
pH
Aqueous SO 2 solutions are acidic.
Chemical properties
SO 2 is a powerful reducing and oxidizing agent. In the presence of air and a catalyst it oxidizes to SO 3.
SO 2 + O 2 → SO 3
The lone pairs of electrons sometimes make it behave like a Lewis base, in other words, it can react with compounds where there is an atom that is missing electrons.
If SO 2 is in the form of a gas and dry, it does not attack iron, steel, copper-nickel alloys, or nickel-chromium-iron. However, if it is in a liquid or wet state, it causes corrosion to these metals.
Liquid SO 2 with 0.2% water or more produces strong corrosion to iron, brass and copper. It is corrosive to aluminum.
When liquid, it can also attack some plastics, rubbers, and coatings.
Aqueous SO solutions
SO 2 is very soluble in water. It was considered for a long time that in water it forms sulfurous acid H 2 SO 3, but the existence of this acid has not been demonstrated.
In solutions of SO 2 in water the following equilibria occur:
SO 2 + H 2 O ⇔ SO 2.H 2 O
SO 2.H 2 O ⇔ HSO 3 - + H 3 O +
HSO 3 - + H 2 O ⇔ SO 3 2- + H 3 O +
Where HSO 3 - is the bisulfite ion and SO 3 2- is the sulfite ion. The sulfite ion SO 3 2- is produced mainly when an alkali is added to the SO 2 solution.
Aqueous solutions of SO 2 have reducing properties, especially if they are alkaline.
Other properties
- It is extremely stable against heat, even up to 2000 ° C.
- It is not flammable.
Obtaining
SO 2 is obtained by combustion of sulfur (S) in air, although small amounts of SO 3 are also formed.
S + O 2 → SO 2
It can also be produced by heating various sulfides in the air, burning pyrite minerals and minerals containing sulfides, among others.
In the case of iron pyrite, when oxidized, iron oxide (iii) and SO 2 are obtained:
4 FeS 2 + 11 O 2 → 2 Fe 2 O 3 + 8 SO 2 ↑
Presence in nature
SO 2 is released into the atmosphere by the activity of volcanoes (9%) but it is also caused by other natural activities (15%) and by human actions (76%).
Explosive volcanic eruptions cause significant annual fluctuations or variations of SO 2 in the atmosphere. It is estimated that 25% of the SO 2 emitted by volcanoes is washed away by rain before reaching the stratosphere.
Natural sources are the most abundant and are due to the biological cycle of sulfur.
In urban and industrial areas human sources predominate. The main human activity that produces it is the burning of fossil fuels, such as coal, gasoline and diesel. Other human sources are oil refineries, chemical plants, and gas production.
Human activities such as burning coal for electricity are a source of SO 2 pollution. Adrem68. Source: Wikimedia Commons.
In mammals, it is generated endogenously, that is, within the body of animals and humans due to the metabolism of sulfur-containing amino acids (S), especially L-cysteine.
Applications
In the production of sulfuric acid
One of the most important applications of SO 2 is in obtaining sulfuric acid H 2 SO 4.
2 SO 2 + 2 H 2 O + O 2 → 2 H 2 SO 4
In the processed food industry
Sulfur dioxide is used as a food preservative and stabilizer, as a moisture control agent, and as a taste and texture modifier in certain edible products.
It is also used to disinfect equipment that comes into contact with foodstuffs, fermentation equipment, such as those in breweries and wineries, food containers, etc.
It allows you to preserve fruits and vegetables, increases their life on the supermarket shelf, prevents loss of color and flavor and helps in the retention of vitamin C (ascorbic acid) and carotenes (precursors of vitamin A).
Dried fruits are kept free of fungi and bacteria thanks to SO 2. Author: Isabel Ródenas. Source: Pixabay.com
It is used to preserve wine, as it destroys bacteria, fungi and unwanted yeasts. It is also used to sterilize and prevent the formation of nitrosamines in beer.
Barley fermentation equipment to obtain beer is sterilized with SO 2. Author: Cerdadebbie. Source: Pixabay.
It is also used to soak corn kernels, to whiten beet sugar, and as an antimicrobial in the manufacture of high fructose corn syrup.
As a solvent and reagent
It has been widely used as a non-aqueous solvent. Although it is not an ionizing solvent, it is useful as a proton-free solvent for certain analytical applications and chemical reactions.
It is used as a solvent and reagent in organic synthesis, an intermediate in the production of other compounds such as chlorine dioxide, acetyl chloride and in the sulfonation of oils.
As a reducing agent
It is used as a reducing agent despite not being so strong, and in alkaline solution the sulfite ion is formed, which is a more energetic reducing agent.
In various applications
SO 2 is also used:
- In agriculture as a fungicide and preservative for grapes after harvest.
- To manufacture hydrosulfites.
- To bleach wood pulp and paper, as it allows to stabilize the pulp after bleaching with hydrogen peroxide H 2 O 2; SO 2 works by destroying the remaining H 2 O 2 and thus maintain the pulp's brightness, since H 2 O 2 can cause a reversal of brightness.
- To whiten textile fibers and wicker articles.
- To treat water as it eliminates the residual chlorine that remains after the chlorination of drinking water, wastewater or industrial water.
- In the refining of minerals and metals, as a reducing agent for iron during mineral processing.
- In oil refining to trap oxygen and retard corrosion, and as an extraction solvent.
- As an antioxidant.
- As an alkali neutralizer in glass manufacturing.
- In lithium batteries as an oxidizing agent.
Effects of OS
Studies have shown that endogenous or body-produced SO 2 has a beneficial effect on the cardiovascular system, including regulating heart function and relaxing blood vessels.
When SO 2 is produced in the body, it is converted into its derivatives bisulfite HSO 3 - and sulfite SO 3 2-, which exert a vaso-relaxant effect on the arteries.
Endogenous SO 2 reduces hypertension, prevents the development of atherosclerosis, and protects the heart from myocardial damage. It also has an antioxidant action, inhibits inflammation and apoptosis (programmed cell death).
For these reasons it is thought that it may be a possible new therapy for cardiovascular diseases.
The heart can benefit from the SO 2 produced by the body. Author: OpenClipart-Vectors. Source: Pixabay.
Risks
- Exposure to gaseous SO 2 can lead to burns to eyes, skin, throat and mucous membranes, damage to bronchial tubes and lungs.
- Some studies report that it has a potential risk of damage to the genetic material of mammalian and human cells.
- It is corrosive. It is not flammable.
Ecotoxicity
Sulfur dioxide is the most common pollutant gas in the atmosphere, especially in urban and industrial areas.
Its presence in the atmosphere contributes to so-called “acid rain” that is harmful to aquatic organisms, fish, terrestrial vegetation, and corrosion to human-made materials.
Monument damaged by acid rain. Nino Barbieri. Source: Wikimedia Commons.
SO 2 is toxic to fish. Green plants are extremely sensitive to atmospheric SO 2. Alfalfa, cotton, barley, and wheat are damaged at low environmental levels, while potatoes, onions, and corn are much more resistant.
Effects of ingesting it with food
Although it is harmless to healthy people, when used in the concentrations recommended by authorized health agencies, SO 2 can induce asthma in sensitive people who take it with food.
Sensitive people can suffer from asthma when eating foods with small amounts of SO 2. Suraj at Malayalam Wikipedia. Source: Wikimedia Commons.
The foods that usually contain it are dried fruits, artificial soft drinks and alcoholic beverages.
References
- US National Library of Medicine. (2019). Sulfur dioxide. Recovered from pubchem.ncbi.nlm.nih.gov.
- Huang, Y. et al. (2016). Endogenous Sulfur Dioxide: A New Member of Gasotransmitter Family in the Cardiovascular System. Oxid Med Cell Longev. 2016; 2016: 8961951. Recovered from ncbi.nlm.nih.gov.
- Cotton, F. Albert and Wilkinson, Geoffrey. (1980). Advanced Inorganic Chemistry. Fourth Edition. John Wiley & Sons.
- Windholz, M. et al. (editors) (1983). The Merck Index. An Encyclopedia of Chemicals, Drugs, and Biologicals. Tenth Edition. Merck & CO., Inc.
- Pan, X. (2011). Sulfur Oxides: Sources, Exposures and Health Effects. Health Effects of Sulfur Oxides. In Encyclopedia of Environmental Health. Recovered from sciencedirect.com.
- Tricker, R. and Tricker, S. (1999). Pollutants and contaminants. Sulfur dioxide. In Environmental Requirements for Electromechanical and Electronic Equipment. Recovered from sciencedirect.com.
- Bleam, W. (2017). Acid-Base Chemistry. Sulfur Oxides. In Soil and Environmental Chemistry (Second Edition). Recovered from sciencedirect.com.
- Freedman, BJ (1980). Sulfur dioxide in foods and beverages: its use as a preservative and its effect on asthma. Br J Dis Chest. 1980; 14 (2): 128-34. Recovered from ncbi.nlm.nih.gov.
- Craig, K. (2018). A Review of the Chemistry, Pesticide Use, and Environmental Fate of Sulfur Dioxide, as Used in California. In Reviews of Environmental Contamination and Toxicology. Volume 246. Recovered from link.springer.com.