Sulfur dioxide (IUPAC-recommended spelling) or sulphur dioxide (traditional Commonwealth English) is the chemical compound with the formula . It is a colorless gas with a pungent smell that is responsible for the odor of burnt matches. It is released naturally by volcanic activity and is produced anthropogenically as a by-product of metals refining and the burning of fossil fuels for energy.
Sulfur dioxide is somewhat toxic to humans, although only when inhaled in relatively large quantities for a period of several minutes or more. It was known to medieval alchemists as "volatile spirit of sulfur".
Structure and bonding SO2 is a bent molecule with C2v symmetry point group. A valence bond theory approach considering just s and p orbitals would describe the bonding in terms of resonance between two resonance structures.
The sulfur–oxygen bond has a bond order of 1.5. There is support for this simple approach that does not invoke d orbital participation. In terms of electron-counting formalism, the sulfur atom has an oxidation state of +4 and a formal charge of +1.
Occurrence
Sulfur dioxide is found on Earth and exists in very small concentrations in the atmosphere at about 15 ppb.
On other planets, sulfur dioxide can be found in various concentrations, the most significant being the atmosphere of Venus, where it is the third-most abundant atmospheric gas at 150 ppm. There, it reacts with water to form clouds of sulfurous acid ( + ⇌ + ), and is a key component of the planet's global atmospheric sulfur cycle. It has been implicated as a key agent in the warming of early Mars, with estimates of concentrations in the lower atmosphere as high as 100 ppm, though it only exists in trace amounts. On both Venus and Mars, as on Earth, its primary source is thought to be volcanic. The atmosphere of Io, a natural satellite of Jupiter, is 90% sulfur dioxide and trace amounts are thought to also exist in the atmosphere of Jupiter. The James Webb Space Telescope has observed the presence of sulfur dioxide on the exoplanet WASP-39b, where it is formed through photochemistry in the planet's atmosphere.
As an ice, it is thought to exist in abundance on the Galilean moons—as subliming ice or frost on the trailing hemisphere of Io, and in the crust and mantle of Europa, Ganymede, and Callisto, possibly also in liquid form and readily reacting with water.Europa's Hidden Ice Chemistry – NASA Jet Propulsion Laboratory. Jpl.nasa.gov (October 4, 2010). Retrieved on September 24, 2013.
Production Sulfur dioxide is primarily produced for sulfuric acid manufacture (see contact process, but other processes predated that at least since 16th century). In the United States in 1979, 23.6 million metric tons (26 million U.S. short tons) of sulfur dioxide were used in this way, compared with 150,000 metric tons (165,347 U.S. short tons) used for other purposes. Most sulfur dioxide is produced by the combustion of elemental sulfur. Some sulfur dioxide is also produced by roasting pyrite and other sulfide ores in air.
Combustion routes Sulfur dioxide is the product of the burning of sulfur or of burning materials that contain sulfur:
+ 8 → 8 , ΔH = −297 kJ/mol
To aid combustion, liquified sulfur ( is sprayed through an atomizing nozzle to generate fine drops of sulfur with a large surface area. The reaction is exothermic, and the combustion produces temperatures of . The significant amount of heat produced is recovered by steam generation that can subsequently be converted to electricity.
The combustion of hydrogen sulfide and organosulfur compounds proceeds similarly. For example:
2 + 3 → 2 + 2
The roasting of sulfide ores such as pyrite, sphalerite, and cinnabar (mercury sulfide) also releases SO2:Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company; New York, 2010; p. 414.
4 + 11 → 2 + 8 2 + 3 → 2 + 2
4 FeS + 7 → 2 + 4
A combination of these reactions is responsible for the largest source of sulfur dioxide, volcanic eruptions. These events can release millions of tons of SO2.
Reduction of higher oxides Sulfur dioxide can also be a byproduct in the manufacture of calcium silicate cement; CaSO4 is heated with coke and sand in this process:
2 + 2 + C → 2 + 2 +
Until the 1970s commercial quantities of sulfuric acid and cement were produced by this process in Whitehaven, England. Upon being mixed with shale or marl, and roasted, the sulfate liberated sulfur dioxide gas, used in sulfuric acid production, the reaction also produced calcium silicate, a precursor in cement production.WHITEHAVEN COAST ARCHAEOLOGICAL SURVEY. lakestay.co.uk (2007)
On a laboratory scale, the action of hot concentrated sulfuric acid on copper turnings produces sulfur dioxide.
Cu + 2 → Tin also reacts with concentrated sulfuric acid but it produces tin(II) sulfate which can later be pyrolyzed at 360 °C into tin dioxide and dry sulfur dioxide.
Sn + → → +
From sulfites The reverse reaction occurs upon acidification:
Reactions Sulfites result by the action of aqueous base on sulfur dioxide:
Sulfur dioxide is a mild but useful reducing agent. It is oxidized by halogens to give the sulfuryl halides, such as sulfuryl chloride:
Sulfur dioxide is the oxidising agent in the Claus process, which is conducted on a large scale in oil refineries. Here, sulfur dioxide is reduced by hydrogen sulfide to give elemental sulfur:
The sequential oxidation of sulfur dioxide followed by its hydration is used in the production of sulfuric acid. 2 + 2 + → 2 Sulfur dioxide dissolves in water to give "sulfurous acid", which cannot be isolated and is instead an acidic solution of bisulfite, and possibly sulfite, ions. Ka = 1.54; pKa = 1.81
Laboratory reactions Sulfur dioxide is one of the few common acidic yet reducing gases. It turns moist litmus pink (being acidic), then white (due to its bleaching effect). It may be identified by bubbling it through a dichromate solution, turning the solution from orange to green (Cr3+ (aq)). It can also reduce ferric ions to ferrous.
Sulfur dioxide can react with certain 1,3-dienes in a cheletropic reaction to form cyclic sulfones. This reaction is exploited on an industrial scale for the synthesis of sulfolane, which is an important solvent in the petrochemical industry.
Sulfur dioxide can bind to metal ions as a ligand to form metal sulfur dioxide complexes, typically where the transition metal is in oxidation state 0 or +1. Many different bonding modes (geometries) are recognized, but in most cases, the ligand is monodentate, attached to the metal through sulfur, which can be either planar and pyramidal η1. As a η1-SO2 (S-bonded planar) ligand sulfur dioxide functions as a Lewis base using the lone pair on S. SO2 functions as a Lewis acids in its η1-SO2 (S-bonded pyramidal) bonding mode with metals and in its 1:1 adducts with Lewis bases such as dimethylacetamide and trimethyl amine. When bonding to Lewis bases the acid parameters of SO2 are EA = 0.51 and EA = 1.56.
Uses The overarching, dominant use of sulfur dioxide is in the production of sulfuric acid.
Precursor to sulfuric acid Sulfur dioxide is an intermediate in the production of sulfuric acid, being converted to sulfur trioxide, and then to oleum, which is made into sulfuric acid. Sulfur dioxide for this purpose is made when sulfur combines with oxygen. The method of converting sulfur dioxide to sulfuric acid is called the contact process. Several million tons are produced annually for this purpose.
Food preservative
Sulfur dioxide is sometimes used as a preservative for dried apricots, dried figs, and other dried fruits, owing to its antimicrobial properties and ability to prevent oxidation, and is called E220Current EU approved additives and their E Numbers, The Food Standards Agency website. when used in this way in Europe. As a preservative, it maintains the colorful appearance of the fruit and prevents rotting. Historically, molasses was "sulfured" as a preservative and also to lighten its color. Treatment of dried fruit was usually done outdoors, by igniting sublimed sulfur and burning in an enclosed space with the fruits. Fruits may be sulfured by dipping them into sodium bisulfite, sodium sulfite or sodium metabisulfite.
Winemaking Sulfur dioxide was first used in winemaking by the Romans, when they discovered that burning sulfur candles inside empty wine vessels keeps them fresh and free from vinegar smell.
It is still an important compound in winemaking, and is measured in parts per million (ppm) in wine. It is present even in so-called unsulfurated wine at concentrations of up to 10 mg/L.Sulphites in wine, MoreThanOrganic.com. It serves as an antibiotic and antioxidant, protecting wine from spoilage by bacteria and oxidation – a phenomenon that leads to the browning of the wine and a loss of cultivar specific flavors.Jackson, R.S. (2008) Wine science: principles and applications, Amsterdam; Boston: Elsevier/Academic Press Its antimicrobial action also helps minimize volatile acidity. Wines containing sulfur dioxide are typically labeled with "containing sulfites".
Sulfur dioxide exists in wine in free and bound forms, and the combinations are referred to as total SO2. Binding, for instance to the carbonyl group of acetaldehyde, varies with the wine in question. The free form exists in equilibrium between molecular SO2 (as a dissolved gas) and bisulfite ion, which is in turn in equilibrium with sulfite ion. These equilibria depend on the pH of the wine. Lower pH shifts the equilibrium towards molecular (gaseous) SO2, which is the active form, while at higher pH more SO2 is found in the inactive sulfite and bisulfite forms. The molecular SO2 is active as an antimicrobial and antioxidant, and this is also the form which may be perceived as a pungent odor at high levels. Wines with total SO2 concentrations below 10 ppm do not require "contains sulfites" on the label by US and EU laws. The upper limit of total SO2 allowed in wine in the US is 350 ppm; in the EU it is 160 ppm for red wines and 210 ppm for white and rosé wines. In low concentrations, SO2 is mostly undetectable in wine, but at free SO2 concentrations over 50 ppm, SO2 becomes evident in the smell and taste of wine.
SO2 is also a very important compound in winery sanitation. Wineries and equipment must be kept clean, and because bleach cannot be used in a winery due to the risk of cork taint,Chlorine Use in the Winery. Purdue University a mixture of SO2, water, and citric acid is commonly used to clean and sanitize equipment. Ozone (O3) is now used extensively for sanitizing in wineries due to its efficacy, and because it does not affect the wine or most equipment.Use of ozone for winery and environmental sanitation , Practical Winery & Vineyard Journal.
As a soaking agent Aqueous sulfur dioxide solution is used in corn wet-milling, in the steeping stage of the process. The corn kernels are soaked in this solution in large tanks containing lactic acid and sulfur dioxide at around 53˚C (127˚F) temperature for nearly 40 hours. This is done to soften the kernel so that the oil in the germ will not contaminate other products and is easy to separate.
As a reducing agent Sulfur dioxide is also a good reductant. In the presence of water, sulfur dioxide is able to decolorize substances. Specifically, it is a useful reducing bleach for papers and delicate materials such as clothes. This bleaching effect normally does not last very long. Oxygen in the atmosphere reoxidizes the reduced dyes, restoring the color. In municipal wastewater treatment, sulfur dioxide is used to treat chlorinated wastewater prior to release. Sulfur dioxide reduces free and combined chlorine to chloride.