Sulfur dioxide, also known as sulphur dioxide in some regions, is a chemical compound with the formula SO₂. It is a colorless gas with a strong, sharp smell, similar to the scent of burnt matches. This gas is naturally released during volcanic eruptions and is also created as a by-product during processes like refining metals and burning fuels that contain sulfur.
Sulfur dioxide can be harmful to humans if inhaled in large amounts over several minutes or longer. In the past, this gas was called "volatile spirit of sulfur" by medieval alchemists.
Structure and bonding
SO₂ is a molecule with a bent shape and a point group symmetry called C₂v. Using a model that only considers s and p orbitals, the bonding can be explained by showing two different ways that electrons are shared between atoms.
The bond between sulfur and oxygen has a bond order of 1.5, meaning it is stronger than a single bond but weaker than a double bond. This simple explanation does not require considering other types of orbitals. According to electron-counting rules, the sulfur atom has an oxidation state of +4 and a formal charge of +1.
Occurrence
Sulfur dioxide is present on Earth and exists in very small amounts in the atmosphere, about 15 parts per billion (ppb).
On other planets, sulfur dioxide can be found in different amounts. On Venus, it is the third-most common gas in the atmosphere, making up 150 parts per million (ppm). There, sulfur dioxide reacts with water to form clouds of sulfurous acid (SO₂ + H₂O ⇌ HSO₃⁻ + H⁺) and plays an important role in Venus's atmospheric sulfur cycle. On Mars, sulfur dioxide may have contributed to warming the planet in the past, with estimates of up to 100 ppm in the lower atmosphere, though it is now only present in very small amounts. On both Venus and Mars, volcanic activity is believed to be the main source of sulfur dioxide. The moon Io, which orbits Jupiter, has an atmosphere that is 90% sulfur dioxide, and small amounts may also exist in Jupiter's atmosphere. The James Webb Space Telescope has detected sulfur dioxide on the exoplanet WASP-39b, where it forms through chemical reactions caused by sunlight.
As ice, sulfur dioxide is thought to be widespread on the Galilean moons—present as frost or ice that turns directly into gas on the trailing side of Io, and possibly as liquid or solid in the crust and mantle of Europa, Ganymede, and Callisto. It may also react with water in these locations.
Production
Sulfur dioxide is mainly made for producing sulfuric acid, a process known as the contact process. Other methods existed before this, dating back at least to the 16th century. In 1979, the United States used 23.6 million metric tons (26 million U.S. short tons) of sulfur dioxide for making sulfuric acid, compared to 150,000 metric tons (165,347 U.S. short tons) used for other purposes. Most sulfur dioxide is created by burning elemental sulfur. Some is also made by roasting pyrite and other sulfide ores in air.
Sulfur dioxide is formed when sulfur or sulfur-containing materials are burned. To help the burning process, liquid sulfur (140–150 °C or 284–302 °F) is sprayed through a nozzle to create small sulfur droplets with a large surface area. Burning sulfur releases a lot of heat, reaching temperatures of 1,000–1,600 °C (1,830–2,910 °F). The heat is captured to make steam, which can later be used to generate electricity.
Burning hydrogen sulfide and other sulfur-containing compounds works in a similar way. For example, roasting sulfide ores like pyrite, sphalerite, and cinnabar (mercury sulfide) also releases sulfur dioxide.
The largest natural source of sulfur dioxide is volcanic eruptions, which can release millions of tons of the gas.
Sulfur dioxide is also made during the production of calcium silicate cement. In this process, calcium sulfate (CaSO₄) is heated with coke and sand.
Until the 1970s, sulfuric acid and cement were made using this method in Whitehaven, England. When mixed with shale or marl and roasted, the sulfate in the mixture releases sulfur dioxide gas, which is used to make sulfuric acid. This reaction also creates calcium silicate, a material used in cement production.
In laboratories, heating concentrated sulfuric acid on copper produces sulfur dioxide. Tin reacts with concentrated sulfuric acid to form tin(II) sulfate, which can later be heated at 360 °C to produce tin dioxide and sulfur dioxide.
When sulfur dioxide is exposed to acid, the reverse reaction can occur.
Reactions
Sulfites are formed when sulfur dioxide reacts with a water-based solution.
Sulfur dioxide is a mild reducing agent. It can be oxidized by halogens to create sulfuryl halides, such as sulfuryl chloride.
In the Claus process, which is widely used in oil refineries, sulfur dioxide acts as an oxidizing agent. It reacts with hydrogen sulfide to produce elemental sulfur.
The oxidation of sulfur dioxide, followed by its reaction with water, is used to make sulfuric acid.
When sulfur dioxide dissolves in water, it forms "sulfurous acid." This substance cannot be isolated as a pure compound. Instead, it exists as a solution containing bisulfite and possibly sulfite ions.
Sulfur dioxide is a rare example of a gas that is both acidic and reducing. It turns moist litmus paper pink (due to its acidity) and then white (because of its bleaching effect). It can be identified by bubbling it through a dichromate solution, which changes from orange to green (Cr³⁺ ions). It also reduces ferric ions to ferrous ions.
Sulfur dioxide can react with certain 1,3-dienes in a specific reaction to form cyclic sulfones. This reaction is used in industry to produce sulfolane, a valuable solvent in the petrochemical industry.
Sulfur dioxide can bind to metal ions to form metal-sulfur dioxide complexes. These complexes often involve transition metals in oxidation states 0 or +1. In most cases, sulfur dioxide attaches to the metal through sulfur, forming a single bond. It can act as a Lewis base when bonded in a planar shape or as a Lewis acid when bonded in a pyramidal shape or when forming adducts with substances like dimethylacetamide or trimethylamine. When bonding to Lewis bases, sulfur dioxide has acid parameters of E_A = 0.51 and E_A = 1.56.
Uses
The main use of sulfur dioxide is in making sulfuric acid. Sulfur dioxide is an intermediate step in the production of sulfuric acid. It is first converted to sulfur trioxide and then to oleum, which is used to make sulfuric acid. Sulfur dioxide for this purpose is created when sulfur combines with oxygen. The method of converting sulfur dioxide to sulfuric acid is called the contact process. Millions of tons are produced each year for this purpose.
Sulfur dioxide is sometimes used as a preservative for dried apricots, dried figs, and other dried fruits. It helps prevent bacteria and mold and stops the fruit from turning brown. In Europe, it is called E 220 when used as a preservative. It keeps the fruit looking colorful and prevents it from rotting. In the past, molasses was treated with sulfur to preserve it and make it lighter in color. Dried fruits were often treated outdoors by burning sulfur in an enclosed space with the fruit. Today, fruits may be treated by dipping them in solutions like sodium bisulfite, sodium sulfite, or sodium metabisulfite.
Sulfur dioxide was first used in winemaking by the Romans. They discovered that burning sulfur candles inside empty wine containers kept the wine fresh and prevented it from smelling like vinegar. Sulfur dioxide is still important in winemaking. It is measured in parts per million (ppm) in wine. Even wines labeled as "unsulfurated" contain up to 10 mg/L of sulfur dioxide. It acts as an antibiotic and antioxidant, protecting wine from bacteria and oxidation, which causes wine to turn brown and lose its unique flavors. It also helps reduce volatile acidity. Wines containing sulfur dioxide are usually labeled with "contains sulfites."
Sulfur dioxide exists in wine in two forms: free and bound. These forms together are called total SO₂. Free sulfur dioxide exists in balance between molecular SO₂ (a dissolved gas) and bisulfite ion, which is also in balance with sulfite ion. These balances depend on the wine's pH. Lower pH increases the amount of active molecular SO₂, which is the form that works as an antimicrobial and antioxidant. At higher pH, more SO₂ is in inactive forms. Molecular SO₂ can cause a strong smell at high levels. Wines with total SO₂ below 10 ppm do not need "contains sulfites" labels in the U.S. and Europe. The maximum allowed total SO₂ in wine is 350 ppm in the U.S., 160 ppm in red wines in the EU, and 210 ppm in white and rosé wines in the EU. At low levels, sulfur dioxide is not easily noticed in wine. However, when free SO₂ exceeds 50 ppm, its smell and taste become noticeable.
Sulfur dioxide is also important for cleaning in wineries. Wineries and equipment must be kept clean, and bleach is not used because it can cause cork taint. Instead, a mixture of sulfur dioxide, water, and citric acid is commonly used for cleaning. Ozone (O₃) is now widely used in wineries for sanitizing because it is effective and does not harm wine or equipment.
Aqueous sulfur dioxide solution is used in corn wet-milling during the steeping stage. Corn kernels are soaked in this solution with lactic acid and sulfur dioxide at about 53°C (127°F) for nearly 40 hours. This softens the kernels, making it easier to separate oil from the germ and preventing it from contaminating other products.
Sulfur dioxide is a good reductant. In the presence of water, it can remove color from substances. It is used as a bleach for paper and delicate materials like clothing. However, this effect does not last long because oxygen in the air can restore the color. In municipal wastewater treatment, sulfur dioxide is used to remove chlorine from water before it is released. It reduces free and combined chlorine to chloride.
Sulfur dioxide dissolves well in water. However, the hypothetical sulfurous acid (H₂SO₃) is not present in significant amounts. Instead, solutions show the presence of hydrogen sulfite ions (HSO₃⁻), which are the actual reducing agents.
In the early 20th century, sulfur dioxide was used in Buenos Aires to kill rats carrying the Yersinia pestis bacterium, which causes bubonic plague. This method was successful and later used in other parts of South America. Sulfur dioxide treatment machines, called Sulfurozador in Buenos Aires, were also used in Rio de Janeiro, New Orleans, and San Francisco for large-scale disinfection campaigns.
Sulfur dioxide or its related form, bisulfite, is produced naturally by certain bacteria, including sulfate-reducing and sulfur-oxidizing bacteria. Its role in mammalian biology is not fully understood. Sulfur dioxide blocks nerve signals from the pulmonary stretch receptors and stops the Hering–Breuer inflation reflex.
Endogenous sulfur dioxide may play a key role in regulating heart and blood vessel function. Problems with sulfur dioxide metabolism may contribute to conditions like high blood pressure, atherosclerosis, pulmonary arterial hypertension, and chest pain. In children with pulmonary arterial hypertension caused by heart defects, levels of homocysteine are higher, and sulfur dioxide levels are lower than in healthy children. These levels are closely linked to the severity of the condition. Scientists believe homocysteine could be a useful marker for disease severity, and sulfur dioxide metabolism may be a target for treatment.
Endogenous sulfur dioxide reduces the growth of smooth muscle cells in blood vessels by lowering MAPK activity and activating adenylyl cyclase and protein kinase A. This is important because smooth muscle cell growth contributes to high blood pressure and atherosclerosis. At low levels, sulfur dioxide causes blood vessels to widen (vasodilation) by acting on the endothelium. At higher levels, it causes vasodilation without involving the endothelium and reduces heart function, lowering blood pressure and heart oxygen use. These effects are linked to ATP-dependent and L-type calcium channels. Sulfur dioxide is also a strong anti-inflammatory, antioxidant, and cell-protecting agent. It lowers blood pressure, prevents thickening of blood vessel walls, and helps control lipid metabolism. It also reduces heart damage caused by excessive stimulation and strengthens the heart's antioxidant defenses.
Sulfur dioxide is a useful solvent for dissolving highly oxidizing salts. It is occasionally used as a source of the sulfonyl group in organic chemistry. It can react with aryl diazonium salts in the presence of cuprous chloride.
Safety
In the United States, the Center for Science in the Public Interest considers sulfur dioxide and sodium bisulfite safe for most people to eat or drink. However, people with asthma may have serious reactions to these preservatives, especially if they consume large amounts. Symptoms of sensitivity to sulfiting agents, such as sulfur dioxide, can include trouble breathing that may be life-threatening. These symptoms can appear within minutes after eating or drinking something containing these agents. Sulfites may also cause problems in people without asthma, such as skin rashes, hives, redness, low blood pressure, stomach pain, and diarrhea. In rare cases, sulfites can lead to a severe allergic reaction called anaphylaxis.
Exposure to sulfur dioxide happens often in daily life, such as from smoke produced by matches, coal, and fuels like bunker fuel. Compared to other chemicals, sulfur dioxide is only slightly harmful and requires very high amounts to cause serious harm. However, because it is found in many places, it is a major air pollutant that can harm human health.
In 2008, the American Conference of Governmental Industrial Hygienists lowered the short-term exposure limit for sulfur dioxide to 0.25 parts per million (ppm). In the United States, the OSHA set the PEL (permissible exposure limit) at 5 ppm (13 mg/m³) as a time-weighted average. Also in the United States, NIOSH set the IDLH (immediately dangerous to life or health) level at 100 ppm. In 2010, the EPA updated the primary SO₂ NAAQS (National Ambient Air Quality Standards) by creating a new one-hour standard of 75 parts per billion (ppb). The EPA removed the two previous standards because they would not offer extra protection for public health compared to the new one-hour standard of 75 ppb.
Environmental role
Major volcanic eruptions greatly affect the amount of sulfate aerosols in the atmosphere during the years they occur. Eruptions that are ranked 4 or higher on the Volcanic Explosivity Index send sulfur dioxide and water vapor directly into the upper atmosphere, where they react to form sulfate aerosol plumes. Volcanic emissions vary in composition and have complex chemical reactions because of the presence of ash and other elements in the plume. Only certain types of volcanoes, called stratovolcanoes, that contain felsic magma are responsible for these plumes. Shield volcanoes, which have mafic magma, do not produce plumes that reach the upper atmosphere.
Before the Industrial Revolution, the dimethyl sulfide pathway was the main source of sulfate aerosols in years without major volcanic activity. According to the IPCC First Assessment Report from 1990, volcanic emissions in the 1980s were about 10 million tons per year, while dimethyl sulfide emissions were about 40 million tons per year. However, by the 1980s, human-caused sulfur emissions had become "at least as large" as all natural sulfur emissions combined. These emissions increased from less than 3 million tons per year in 1860 to 15 million tons in 1900, 40 million tons in 1940, and about 80 million tons in 1980. The report also noted that in industrialized regions of Europe and North America, human-caused emissions were about 10 times greater than natural emissions. In the eastern United States during the early 2000s, sulfate particles made up 25% or more of all air pollution. Exposure to sulfur dioxide from coal power plants was linked to a 2.1 times greater risk of death compared to exposure to other types of fine particle pollution.
In the Southern Hemisphere, sulfur concentrations were much lower due to fewer people living there, with about 90% of the global population in the Northern Hemisphere. In the early 1990s, human-caused sulfur emissions dominated in the Northern Hemisphere, where only 16% of annual sulfur emissions were natural. However, natural emissions made up less than half of the total emissions in the Southern Hemisphere.
The increase in sulfate aerosol emissions had many effects. One of the most visible effects was acid rain, which occurs when clouds carrying sulfate aerosols in the lower atmosphere release acidic precipitation. At its worst, acid rain caused the loss of fish and insect life in lakes and streams in sensitive areas, such as the Adirondack Mountains in the United States. Acid rain harms soil by reducing beneficial microorganisms and releasing harmful metals like aluminum, while washing away essential nutrients like magnesium. Plants that cannot survive in more acidic conditions die, with mountain forests being especially affected because they are regularly exposed to sulfate-rich fog at high altitudes. Although acid rain was not strong enough to directly harm human health, breathing air with high sulfate levels can worsen heart and lung conditions, such as asthma and bronchitis. This pollution is also linked to preterm birth and low birth weight. A study of 74,671 pregnant women in Beijing found that each additional 100 μg/m³ of sulfur dioxide in the air reduced infants’ weight by 7.3 grams, making air pollution the largest risk factor for low birth weight observed.
In the United States, the Acid Rain Program led to a 33% decrease in sulfur emissions between 1983 and 2002. This improvement was partly due to flue-gas desulfurization, a technology that captures sulfur dioxide from power plants burning coal or petroleum. Calcium oxide (lime) reacts with sulfur dioxide to form calcium sulfite. When this compound is exposed to oxygen, it becomes calcium sulfate, or anhydrite. Most gypsum sold in Europe comes from this process.
To control sulfur emissions, many methods have been developed for coal-fired power plants. Sulfur can be removed during burning by using limestone in fluidized bed combustion. Sulfur can also be removed from fuels before burning. The Claus process is used in refineries to produce sulfur as a byproduct. The Stretford process and redox processes using iron oxides, such as Lo-Cat or Sulferox, are also used. Fuel additives like calcium or magnesium carboxylate can be used in marine engines to reduce sulfur dioxide emissions.
Sulfur dioxide in the upper atmosphere can contribute to ozone depletion when combined with chlorofluorocarbons and other halogenated compounds. However, volcanic eruptions can also have complex effects on the ozone layer. For example, the 2022 eruption of Hunga Tonga-Hunga Haʻapai injected sulfur dioxide and water vapor into the stratosphere, altering atmospheric circulation and reducing ozone in the Southern Hemisphere while increasing it in the tropics. The presence of hydrochloric acid in eruptions can also lead to ozone depletion.
Research found that sunlight reaching Earth’s surface decreased by about 4–5% per decade from the late 1950s to the 1980s, and by 2–3% per decade when the 1990s were included. This dimming was not caused by changes in the sun itself, as solar radiation at the top of the atmosphere changed by less than 0.1–0.3% during this time. The dimming affected visible and infrared light but not ultraviolet light. It also occurred on clear days, showing it was not caused by clouds alone.
Changes in aerosol concentrations influence global climate and must be considered when predicting future warming. Models estimate that the cooling effect of sulfates is similar to the warming effect of methane, a short-lived greenhouse gas. However, recent years have seen methane levels rise faster than in the 1980s, driven by wetland emissions, while air pollution has decreased. These trends suggest that 1.5°C warming may occur around 2030 instead of the mid-2010s.
As the importance of sulfate aerosols to the climate became clear, research on their formation and effects increased.