Sulfur dioxide is the chemical compound with the formula SO
2. At standard atmosphere, it is a toxic gas with a pungent, irritating, and rotten smell.
The triple point is 197.69 K and 1.67kPa. It is released naturally by volcanic activity.
Sulfur dioxide was used by the Romans in winemaking, when they discovered that burning sulfur candles
inside empty wine vessels kept them fresh and free from vinegar smell.
Structure and bonding SO2 is a bent molecule with C2v symmetry point
group. A valence bond theory approach which just considered 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 It is found on Earth and exists in very small
concentrations and in the atmosphere at about 1 ppb.
On other planets, it can be found in various concentrations, the most significant being
the atmosphere of Venus, where it is the third-most significant atmospheric gas at 150 ppm. There,
it condenses to form clouds, and is a key component of chemical reactions in the planet's
atmosphere and contributes to global warming. 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, its primary source, like
on Earth, is believed to be volcanic. It is also believed to exist in trace amounts in
the atmospheres of Jupiter and Jupiter. As an ice, it is thought to exist in abundance
on the Galilean moons – as sublimating ice or frost on the trailing hemisphere of Io,
a natural satellite of Jupiter and in the crust and mantle of Europa, Ganymede, and
Callisto, possibly also in liquid form and readily reacting with water.
Production Sulfur dioxide is primarily produced for sulfuric
acid manufacture. In the United States in 1979, 23.6 million tonnes of sulfur dioxide
were used in this way, compared with 150 thousand tonnes 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:
S + O2 → SO2, ΔH = -297 kJ/mol To aid combustion, liquefied 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 1000–1600°C. 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 H2S + 3 O2 → 2 H2O + 2 SO2
The roasting of sulfide ores such as pyrite, sphalerite, and cinnabar also releases SO2:
4 FeS2 + 11 O2 → 2 Fe2O3 + 8 SO2 2 ZnS + 3 O2 → 2 ZnO + 2 SO2
HgS + O2 → Hg + SO2 4 FeS + 7O2 → 2 Fe2O3 + 4 SO2
A combination of these reactions is responsible
for the largest source of sulfur dioxide, volcanic eruptions. These events can release
millions of tonnes 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 CaSO4 + 2 SiO2 + C → 2 CaSiO3 + 2 SO2 + CO2
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.
On a laboratory scale, the action of hot sulfuric acid on copper turnings produces sulfur dioxide.
Cu + 2 H2SO4 → CuSO4 + SO2 + 2 H2O From sulfite
Sulfite results from the reaction of aqueous base and sulfur dioxide. The reverse reaction
involves acidification of sodium metabisulfite: H2SO4 + Na2S2O5 → 2 SO2 + Na2SO4 + H2O
Reactions Industrial reactions
Treatment of basic solutions with sulfur dioxide affords sulfite salts:
SO2 + 2 NaOH → Na2SO3 + H2O Featuring sulfur in the +4 oxidation state,
sulfur dioxide is a reducing agent. It is oxidized by halogens to give the sulfuryl
halides, such as sulfuryl chloride: SO2 + Cl2 → SO2Cl2
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: SO2 + 2 H2S → 3 S + 2 H2O
The sequential oxidation of sulfur dioxide followed by its hydration is used in the production
of sulfuric acid. 2 SO2 + 2 H2O + O2 → 2 H2SO4
Laboratory reactions Sulfur dioxide is one of the few common acidic
yet reducing gases. It turns moist litmus pink, then white. It may be identified by
bubbling it through a dichromate solution, turning the solution from orange to green).
It can also reduce ferric ions to ferrous Sulfur dioxide can react with certain 1,3-dienes
in a cheletropic cycloaddition reaction to give organosulfur compounds.
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 are recognized, but in most cases, the ligand is monodentate, attached to the metal
through sulfur, which can be either planar and pyramidal η1.
Uses 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 billion kilograms are produced annually
for this purpose. As a preservative
Sulfur dioxide is sometimes used as a preservative for dried apricots, dried figs, and other
dried fruits, owing to its antimicrobial properties, and is called E220. when used in this way
in Europe. As a preservative, it maintains the colorful appearance of the fruit and prevents
rotting. It is also added to sulfured molasses. In winemaking
Sulfur dioxide was used by the Romans in winemaking, when they discovered that burning sulfur candles
inside empty wine vessels keeps them fresh and free from vinegar smell.
Sulfur dioxide is still an important compound in winemaking, and is measured in parts per
million in wine It is present even in so-called unsulfurated wine at concentrations of up
to 10 mg/L. It serves as an antibiotic and antioxidant, protecting wine from spoilage
by bacteria and oxidation. Its antimicrobial action also helps to minimize volatile acidity.
Sulfur dioxide is responsible for the words "contains sulfites" found on wine labels.
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 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 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 odour 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 nose
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 the risk of cork taint, a mixture of SO2, water, and citric acid is
commonly used to clean and sanitize equipment. Compounds of ozone are now used extensively
as cleaning products in wineries due to their efficiency, and because these compounds do
not affect the wine or equipment. 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.
Sulfur dioxide is fairly soluble in water, and by both IR and Raman spectroscopy; the
hypothetical sulfurous acid, H2SO3, is not present to any extent. However, such solutions
do show spectra of the hydrogen sulfite ion, HSO3−, by reaction with water, and it is
in fact the actual reducing agent present: SO2 + H2O ⇌ HSO3− + H+
Biochemical and biomedical roles Sulfur dioxide is toxic in large amounts.
It or its conjugate base bisulfite is produced biologically as an intermediate in both sulfate-reducing
organisms and in sulfur-oxidizing bacteria, as well. The role of sulfur dioxide in mammalian
biology is not yet well understood. Sulfur dioxide blocks nerve signals from the pulmonary
stretch receptors and abolishes the Hering–Breuer inflation reflex.
As a refrigerant Being easily condensed and possessing a high
heat of evaporation, sulfur dioxide is a candidate material for refrigerants. Prior to the development
of chlorofluorocarbons, sulfur dioxide was used as a refrigerant in home refrigerators.
As a reagent and solvent in the laboratory Sulfur dioxide is a versatile inert solvent
widely used for dissolving highly oxidizing salts. It is also used occasionally as a source
of the sulfonyl group in organic synthesis. Treatment of aryl diazonium salts with sulfur
dioxide and cuprous chloride yields the corresponding aryl sulfonyl chloride, for example:
As an air pollutant
Sulfur dioxide is a noticeable component in the atmosphere, especially following volcanic
eruptions. According to the United States Environmental Protection Agency, the amount
of sulfur dioxide released in the U.S. per year was:
Sulfur dioxide is a major air pollutant and has significant impacts upon human health.
In addition, the concentration of sulfur dioxide in the atmosphere can influence the habitat
suitability for plant communities, as well as animal life. Sulfur dioxide emissions are
a precursor to acid rain and atmospheric particulates. Due largely to the US EPA’s Acid Rain Program,
the U.S. has had a 33% decrease in emissions between 1983 and 2002. This improvement resulted
in part from flue-gas desulfurization, a technology that enables SO2 to be chemically bound in
power plants burning sulfur-containing coal or oil. In particular, calcium oxide reacts
with sulfur dioxide to form calcium sulfite: CaO + SO2 → CaSO3
Aerobic oxidation of the CaSO3 gives CaSO4, anhydrite. Most gypsum sold in Europe comes
from flue-gas desulfurization. Sulfur can be removed from coal during the
burning process by using limestone as a bed material in fluidized bed combustion.
Sulfur can also be removed from fuels prior to burning the fuel, preventing the formation
of SO2 because no sulfur remains in the fuel from which SO2 can be formed. The Claus process
is used in refineries to produce sulfur as a byproduct. The Stretford process has also
been used to remove sulfur from fuel. Redox processes using iron oxides can also be used,
for example, Lo-Cat or Sulferox. Fuel additives, such as calcium additives
and magnesium oxide, are used in gasoline and diesel engines to lower the emission of
sulfur dioxide gases into the atmosphere. As of 2006, China was the world's largest
sulfur dioxide polluter, with 2005 emissions estimated to be 25,490,000 short tons. This
amount represents a 27% increase since 2000, and is roughly comparable with U.S. emissions
in 1980. Safety
Inhalation Inhaling sulfur dioxide is associated with
increased respiratory symptoms and disease, difficulty in breathing, and premature death.
In 2008, the American Conference of Governmental Industrial Hygienists reduced the short-term
exposure limit to 5 parts per million. The OSHA PEL is currently set at 5 ppm time-weighted
average. NIOSH has set the IDLH at 100 ppm. In 2010, the EPA "revised the primary SO2
NAAQS by establishing a new one-hour standard at a level of 75 parts per billion. EPA revoked
the two existing primary standards because they would not provide additional public health
protection given a one-hour standard at 75 ppb."
A 2011 systematic review concluded that exposure to sulfur dioxide is associated with preterm
birth. Ingestion
In the United States, the Center for Science in the Public Interest lists the two food
preservatives, sulfur dioxide and sodium bisulfite, as being safe for human consumption except
for certain asthmatic individuals who may be sensitive to them, especially in large
amounts. Symptoms of sensitivity to sulfiting agents, including sulfur dioxide, manifest
as potentially life-threatening trouble breathing within minutes of ingestion.
See also Sulfur trioxide
Sulfur-iodine cycle National Ambient Air Quality Standards
References
External links United States Environmental Protection Agency
Sulfur Dioxide page International Chemical Safety Card 0074
IARC Monographs. "Sulfur Dioxide and some Sulfites, Bisulfites and Metabisulfites" v54.
1992. p131. NIOSH Pocket Guide to Chemical Hazards
CDC – Sulfure Dioxide – NIOSH Workplace Safety and Health Topic
Sulfur Dioxide, Molecule of the Month