Hypervalent molecule

From Wikipedia, the free encyclopedia

 

A hypervalent molecule (the phenomenon is sometimes colloquially known as expanded octet) is a molecule that contains one or more main group elements apparently bearing more than eight electrons in their valence shells. Phosphorus pentachloride (PCl₅), sulfur hexafluoride (SF₆), chlorine trifluoride (ClF₃), the chlorite (ClO₂⁻) ion, and the triiodide (I₃⁻) ion are examples of hypervalent molecules.

Definitions and nomenclature

Hypervalent molecules were first formally defined by Jeremy I. Musher in 1969 as molecules having central atoms of group 15–18 in any valence other than the lowest (i.e. 3, 2, 1, 0 for Groups 15, 16, 17, 18 respectively, based on the octet rule).

Several specific classes of hypervalent molecules exist:

• Hypervalent iodine compounds are useful reagents in organic chemistry (e.g. Dess–Martin periodinane)
• Tetra-, penta- and hexavalent phosphorus, silicon, and sulfur compounds (ex. PCl₅, PF₅, SF₆, sulfuranes and persulfuranes)
• Noble gas compounds (ex. xenon tetrafluoride, XeF₄)
• Halogen polyfluorides (ex. ClF₅)

N-X-L notation

N-X-L nomenclature, introduced collaboratively by the research groups of Martin, Arduengo, and Kochi in 1980, is often used to classify hypervalent compounds of main group elements, where:

• N represents the number of valence electrons
• X is the chemical symbol of the central atom
• L the number of ligands to the central atom

Examples of N-X-L nomenclature include:

• XeF₂, 10-Xe-2
• PCl₅, 10-P-5
• SF₆, 12-S-6
• IF₇, 14-I-7

History and controversy

The debate over the nature and classification of hypervalent molecules goes back to Gilbert N. Lewis and Irving Langmuir and the debate over the nature of the chemical bond in the 1920s. Lewis maintained the importance of the two-center two-electron (2c-2e) bond in describing hypervalence, thus using expanded octets to account for such molecules. Using the language of orbital hybridization, the bonds of molecules like PF₅ and SF₆ were said to be constructed from sp³dⁿ orbitals on the central atom. Langmuir, on the other hand, upheld the dominance of the octet rule and preferred the use of ionic bonds to account for hypervalence without violating the rule (e.g. "SF₄²⁺ 2F⁻" for SF₆). 

In the late 1920s and 1930s, Sugden argued for the existence of a two-center one-electron (2c-1e) bond and thus rationalized bonding in hypervalent molecules without the need for expanded octets or ionic bond character; this was poorly accepted at the time. In the 1940s and 1950s, Rundle and Pimentel popularized the idea of the three-center four-electron bond, which is essentially the same concept which Sugden attempted to advance decades earlier; the three-center four-electron bond can be alternatively viewed as consisting of two collinear two-center one-electron bonds, with the remaining two nonbonding electrons localized to the ligands.

The attempt to actually prepare hypervalent organic molecules began with Hermann Staudinger and Georg Wittig in the first half of the twentieth century, who sought to challenge the extant valence theory and successfully prepare nitrogen and phosphorus-centered hypervalent molecules. The theoretical basis for hypervalency was not delineated until J.I. Musher's work in 1969.

In 1990, Magnusson published a seminal work definitively excluding the significance of d-orbital hybridization in the bonding of hypervalent compounds of second-row elements. This had long been a point of contention and confusion in describing these molecules using molecular orbital theory. Part of the confusion here originates from the fact that one must include d-functions in the basis sets used to describe these compounds (or else unreasonably high energies and distorted geometries result), and the contribution of the d-function to the molecular wavefunction is large. These facts were historically interpreted to mean that d-orbitals must be involved in bonding. However, Magnusson concludes in his work that d-orbital involvement is not implicated in hypervalency.

Nevertheless, a 2013 study showed that although the Pimentel ionic model best accounts for the bonding of hypervalent species, the energetic contribution of an expanded octet structure is also not null. In this modern valence bond theory study of the bonding of xenon difluoride, it was found that ionic structures account for about 81% of the overall wavefunction, of which 70% arises from ionic structures employing only the p orbital on xenon while 11% arises from ionic structures employing an ${\displaystyle \mathrm {sd} _{z^{2}}}$hybrid on xenon. The contribution of a formally hypervalent structure employing an orbital of sp³d hydridization on xenon accounts for 11% of the wavefunction, with a diradical contribution making up the remaining 8%. The 11% sp³d contribution results in a net stabilization of the molecule by 7.2 kcal (30 kJ) mol⁻¹, a minor but significant fraction of the total energy of the total bond energy (64 kcal (270 kJ) mol⁻¹). Other studies have similarly found minor but non-negligible energetic contributions from expanded octet structures in SF₆ (17%) and XeF₆ (14%).

Despite the lack of chemical realism, the IUPAC recommends the drawing of expanded octet structures for functional groups like sulfones and phosphoranes, in order to avoid the drawing of a large number of formal charges or partial single bonds.

Criticism

Both the term and concept of hypervalency still fall under criticism. In 1984, in response to this general controversy, Paul von Ragué Schleyer proposed the replacement of 'hypervalency' with use of the term hypercoordination because this term does not imply any mode of chemical bonding and the question could thus be avoided altogether.

The concept itself has been criticized by Ronald Gillespie who, based on an analysis of electron localization functions, wrote in 2002 that "as there is no fundamental difference between the bonds in hypervalent and non-hypervalent (Lewis octet) molecules there is no reason to continue to use the term hypervalent."

For hypercoordinated molecules with electronegative ligands such as PF₅ it has been demonstrated that the ligands can pull away enough electron density from the central atom so that its net content is again 8 electrons or fewer. Consistent with this alternative view is the finding that hypercoordinated molecules based on fluorine ligands, for example PF₅ do not have hydride counterparts e.g. phosphorane PH₅ which is unknown. 

The ionic model holds up well in thermochemical calculations. It predicts favorable exothermic formation of PF₄⁺F⁻ from phosphorus trifluoride PF₃ and fluorine F₂ whereas a similar reaction forming PH₄⁺H⁻ is not favorable.

Alternative definition

Durrant has proposed an alternative definition of hypervalency, based on the analysis of atomic charge maps obtained from Atoms in molecules theory. This approach defines a parameter called the valence electron equivalent, γ, as “the formal shared electron count at a given atom, obtained by any combination of valid ionic and covalent resonance forms that reproduces the observed charge distribution”. For any particular atom X, if the value of γ(X) is greater than 8, that atom is hypervalent. Using this alternative definition, many species such as PCl₅, SO₄²⁻, and XeF₄, that are hypervalent by Musher's definition, are reclassified as hypercoordinate but not hypervalent, due to strongly ionic bonding that draws electrons away from the central atom. On the other hand, some compounds that are normally written with ionic bonds in order to conform to the octet rule, such as ozone O₃, nitrous oxide NNO, and trimethylamine N-oxide (CH₃)₃NO, are found to be genuinely hypervalent. Examples of γ calculations for phosphate PO₄³⁻ (γ(P) = 2.6, non-hypervalent) and orthonitrate NO₄³⁻ (γ(N) = 8.5, hypervalent) are shown below.

Calculation of the valence electron equivalent for phosphate and orthonitrate

Bonding in hypervalent molecules

Early considerations of the geometry of hypervalent molecules returned familiar arrangements that were well explained by the VSEPR model for atomic bonding. Accordingly, AB₅ and AB₆ type molecules would possess a trigonal bi-pyramidal and octahedral geometry, respectively. However, in order to account for the observed bond angles, bond lengths and apparent violation of the Lewis octet rule, several alternative models have been proposed. 

In the 1950s an expanded valence shell treatment of hypervalent bonding was adduced to explain the molecular architecture, where the central atom of penta- and hexacoordinated molecules would utilize d AOs in addition to s and p AOs. However, advances in the study of ab initio calculations have revealed that the contribution of d-orbitals to hypervalent bonding is too small to describe the bonding properties, and this description is now regarded as much less important. It was shown that in the case of hexacoordinated SF₆, d-orbitals are not involved in S-F bond formation, but charge transfer between the sulfur and fluorine atoms and the apposite resonance structures were able to account for the hypervalency.

Additional modifications to the octet rule have been attempted to involve ionic characteristics in hypervalent bonding. As one of these modifications, in 1951, the concept of the 3-center 4-electron (3c-4e) bond, which described hypervalent bonding with a qualitative molecular orbital, was proposed. The 3c-4e bond is described as three molecular orbitals given by the combination of a p atomic orbital on the central atom and an atomic orbital from each of the two ligands on opposite sides of the central atom. Only one of the two pairs of electrons is occupying a molecular orbital that involves bonding to the central atom, the second pair being non-bonding and occupying a molecular orbital composed of only atomic orbitals from the two ligands. This model in which the octet rule is preserved was also advocated by Musher.

Qualitative model for a three-center four-electron bond

Molecular orbital theory

A complete description of hypervalent molecules arises from consideration of molecular orbital theory through quantum mechanical methods. A LCAO in, for example, sulfur hexafluoride, taking a basis set of the one sulfur 3s-orbital, the three sulfur 3p-orbitals, and six octahedral geometry symmetry-adapted linear combinations (SALCs) of fluorine orbitals, a total of ten molecular orbitals are obtained (four fully occupied bonding MOs of the lowest energy, two fully occupied intermediate energy non-bonding MOs and four vacant antibonding MOs with the highest energy) providing room for all 12 valence electrons. This is a stable configuration only for SX₆ molecules containing electronegative ligand atoms like fluorine, which explains why SH₆ is not a stable molecule. In the bonding model, the two non-bonding MOs (1e_g) are localized equally on all six fluorine atoms.

Valence bond theory

For hypervalent compounds in which the ligands are more electronegative than the central, hypervalent atom, resonance structures can be drawn with no more than four covalent electron pair bonds and completed with ionic bonds to obey the octet rule. For example, in phosphorus pentafluoride (PF₅), 5 resonance structures can be generated each with four covalent bonds and one ionic bond with greater weight in the structures placing ionic character in the axial bonds, thus satisfying the octet rule and explaining both the observed trigonal bipyramidal molecular geometry and the fact that the axial bond length (158 pm) is longer than the equatorial (154 pm).

Phosphorus pentafluoride. There are 2 structures with an axial ionic bond, plus 3 structures with an equatorial ionic bond.

 

For a hexacoordinate molecule such as sulfur hexafluoride, each of the six bonds is the same length. The rationalization described above can be applied to generate 15 resonance structures each with four covalent bonds and two ionic bonds, such that the ionic character is distributed equally across each of the sulfur-fluorine bonds. 

Sulfur hexafluoride. There are 12 structures with the two ionic bonds in adjacent (cis) positions, plus 3 structures with the two ionic bonds in opposite (trans) positions.

 

Spin-coupled valence bond theory has been applied to diazomethane and the resulting orbital analysis was interpreted in terms of a chemical structure in which the central nitrogen has five covalent bonds.

Chemical formula of diazomethane, showing hypervalent nitrogen

 

This led the authors to the interesting conclusion that "Contrary to what we were all taught as undergraduates, the nitrogen atom does indeed form five covalent linkages and the availability or otherwise of d-orbitals has nothing to do with this state of affairs."

Structure, reactivity, and kinetics

Structure

Hexacoordinated phosphorus

Hexacoordinate phosphorus molecules involving nitrogen, oxygen, or sulfur ligands provide examples of Lewis acid-Lewis base hexacoordination. For the two similar complexes shown below, the length of the C-P bond increases with decreasing length of the N-P bond; the strength of the C-P bond decreases with increasing strength of the N-P Lewis acid-Lewis base interaction. 

Relative bond strengths in hexacoordinated phosphorus compounds. In A, the N-P bond is 1.980Å long and the C-P is 1.833Å long, and in B, the N-P bond increases to 2.013Å as the C-P bond decreases to 1.814Å.

Pentacoordinated silicon

This trend is also generally true of pentacoordinated main-group elements with one or more lone-pair-containing ligand, including the oxygen-pentacoordinated silicon examples shown below. 

Relative bond strengths in pentacoordinated silicon compounds. In A, the Si-O bond length is 1.749Å and the Si-I bond length is 3.734Å; in B, the Si-O bond lengthens to 1.800Å and the Si-Br bond shortens to 3.122Å, and in C, the Si-O bond is the longest at 1.954Å and the Si-Cl bond the shortest at 2.307A.

 

The Si-halogen bonds range from close to the expected van der Waals value in A (a weak bond) almost to the expected covalent single bond value in C (a strong bond).

Reactivity

Silicon

Observed third-order reaction rate constants
for hydrolysis (displacement of chloride from silicon)

Chlorosilane	Nucleophile	kobs (M−2s−1) at 20 °C in anisole
Ph3SiCl	HMPT	1200
Ph3SiCl	DMSO	50
Ph3SiCl	DMF	6
MePh2SiCl	HMPT	2000
MePh2SiCl	DMSO	360
MePh2SiCl	DMF	80
Me(1-Np)PhSiCl	HMPT	3500
Me(1-Np)PhSiCl	DMSO	180
Me(1-Np)PhSiCl	DMF	40
(1-Np)Ph(vinyl)SiCl	HMPT	2200
(1-Np)Ph(vinyl)SiCl	DMSO	90
(1-Np)(m-CF3Ph)HSiCl	DMSO	1800
(1-Np)(m-CF3Ph)HSiCl	DMF	300

Corriu and coworkers performed early work characterizing reactions thought to proceed through a hypervalent transition state. Measurements of the reaction rates of hydrolysis of tetravalent chlorosilanes incubated with catalytic amounts of water returned a rate that is first order in chlorosilane and second order in water. This indicated that two water molecules interacted with the silane during hydrolysis and from this a binucleophilic reaction mechanism was proposed. Corriu and coworkers then measured the rates of hydrolysis in the presence of nucleophilic catalyst HMPT, DMSO or DMF. It was shown that the rate of hydrolysis was again first order in chlorosilane, first order in catalyst and now first order in water. Appropriately, the rates of hydrolysis also exhibited a dependence on the magnitude of charge on the oxygen of the nucleophile. 

Taken together this led the group to propose a reaction mechanism in which there is a pre-rate determining nucleophilic attack of the tetracoordinated silane by the nucleophile (or water) in which a hypervalent pentacoordinated silane is formed. This is followed by a nucleophilic attack of the intermediate by water in a rate determining step leading to hexacoordinated species that quickly decomposes giving the hydroxysilane.

Silane hydrolysis was further investigated by Holmes and coworkers in which tetracoordinated Mes₂SiF₂ (Mes = mesityl) and pentacoordinated Mes₂SiF₃⁻ were reacted with two equivalents of water. Following twenty-four hours, almost no hydrolysis of the tetracoordinated silane was observed, while the pentacoordinated silane was completely hydrolyzed after fifteen minutes. Additionally, X-ray diffraction data collected for the tetraethylammonium salts of the fluorosilanes showed the formation of hydrogen bisilonate lattice supporting a hexacoordinated intermediate from which HF₂⁻ is quickly displaced leading to the hydroxylated product. This reaction and crystallographic data support the mechanism proposed by Corriu et al.

 

Mechanism of silane hydrolysis and structure of the hydrogen bisilonate lattice

 

The apparent increased reactivity of hypervalent molecules, contrasted with tetravalent analogues, has also been observed for Grignard reactions. The Corriu group measured Grignard reaction half-times by NMR for related 18-crown-6 potassium salts of a variety of tetra- and pentacoordinated fluorosilanes in the presence of catalytic amounts of nucleophile.

Though the half reaction method is imprecise, the magnitudinal differences in reactions rates allowed for a proposed reaction scheme wherein, a pre-rate determining attack of the tetravalent silane by the nucleophile results in an equilibrium between the neutral tetracoordinated species and the anionic pentavalent compound. This is followed by nucleophilic coordination by two Grignard reagents as normally seen, forming a hexacoordinated transition state and yielding the expected product.

Grignard reaction mechanism for tetracoordinate silanes and the analogous hypervalent pentacoordinated silanes

 

The mechanistic implications of this are extended to a hexacoordinated silicon species that is thought to be active as a transition state in some reactions. The reaction of allyl- or crotyl-trifluorosilanes with aldehydes and ketones only precedes with fluoride activation to give a pentacoordinated silicon. This intermediate then acts as a Lewis acid to coordinate with the carbonyl oxygen atom. The further weakening of the silicon–carbon bond as the silicon becomes hexacoordinate helps drive this reaction.

Phosphorus

Similar reactivity has also been observed for other hypervalent structures such as the miscellany of phosphorus compounds, for which hexacoordinated transition states have been proposed. Hydrolysis of phosphoranes and oxyphosphoranes have been studied  and shown to be second order in water. Bel'skii et al.. have proposed a prerate determining nucleophilic attack by water resulting in an equilibrium between the penta- and hexacoordinated phosphorus species, which is followed by a proton transfer involving the second water molecule in a rate determining ring-opening step, leading to the hydroxlyated product. 

Mechanism of the hydrolysis of pentacoordinated phosphorus

 

Alcoholysis of pentacoordinated phosphorus compounds, such as trimethoxyphospholene with benzyl alcohol, have also been postulated to occur through a similar octahedral transition state, as in hydrolysis, however without ring opening.

Mechanism of the base catalyzed alcoholysis of pentacoordinated phosphorus

 

It can be understood from these experiments that the increased reactivity observed for hypervalent molecules, contrasted with analogous nonhypervalent compounds, can be attributed to the congruence of these species to the hypercoordinated activated states normally formed during the course of the reaction.

Ab initio calculations

The enhanced reactivity at pentacoordinated silicon is not fully understood. Corriu and coworkers suggested that greater electropositive character at the pentavalent silicon atom may be responsible for its increased reactivity. Preliminary ab initio calculations supported this hypothesis to some degree, but used a small basis set.

A software program for ab initio calculations, Gaussian 86, was used by Dieters and coworkers to compare tetracoordinated silicon and phosphorus to their pentacoordinate analogues. This ab initio approach is used as a supplement to determine why reactivity improves in nucleophilic reactions with pentacoordinated compounds. For silicon, the 6-31+G* basis set was used because of its pentacoordinated anionic character and for phosphorus, the 6-31G* basis set was used.

Pentacoordinated compounds should theoretically be less electrophilic than tetracoordinated analogues due to steric hindrance and greater electron density from the ligands, yet experimentally show greater reactivity with nucleophiles than their tetracoordinated analogues. Advanced ab initio calculations were performed on series of tetracoordinated and pentacoordinated species to further understand this reactivity phenomenon. Each series varied by degree of fluorination. Bond lengths and charge densities are shown as functions of how many hydride ligands are on the central atoms. For every new hydride, there is one less fluoride.

For silicon and phosphorus bond lengths, charge densities, and Mulliken bond overlap, populations were calculated for tetra and pentacoordinated species by this ab initio approach. Addition of a fluoride ion to tetracoordinated silicon shows an overall average increase of 0.1 electron charge, which is considered insignificant. In general, bond lengths in trigonal bipyramidal pentacoordinate species are longer than those in tetracoordinate analogues. Si-F bonds and Si-H bonds both increase in length upon pentacoordination and related effects are seen in phosphorus species, but to a lesser degree. The reason for the greater magnitude in bond length change for silicon species over phosphorus species is the increased effective nuclear charge at phosphorus. Therefore, silicon is concluded to be more loosely bound to its ligands.

Effects of fluorine substitution on positive charge density:

 

Comparison of Charge Densities with Degree of Fluorination for Tetra and Pentacoordinated Silicon

 

In addition Dieters and coworkers show an inverse correlation between bond length and bond overlap for all series. Pentacoordinated species are concluded to be more reactive because of their looser bonds as trigonal-bipyramidal structures.

Calculated bond length and bond overlap with degree of fluorination:

 

Comparison of Bond Lengths with Degree of Fluorination for Tetra and Pentacoordinated Silicon

 

Comparison of Bond Lengths with Degree of Fluorination for Tetra and Pentacoordinated Phosphorus

 

By calculating the energies for the addition and removal of a fluoride ion in various silicon and phosphorus species, several trends were found. In particular, the tetracoordinated species have much higher energy requirements for ligand removal than do pentacoordinated species. Further, silicon species have lower energy requirements for ligand removal than do phosphorus species, which is an indication of weaker bonds in silicon.

Sulfur dioxide

From Wikipedia, the free encyclopedia

 

Sulfur dioxide

Names
IUPAC name Sulfur dioxide
Other names Sulfurous anhydride Sulfur(IV) oxide
Identifiers
CAS Number	7446-09-5
3D model (JSmol)	Interactive image
Beilstein Reference	3535237
ChEBI	CHEBI:18422
ChEMBL	ChEMBL1235997
ChemSpider	1087
ECHA InfoCard	100.028.359
EC Number	231-195-2
E number	E220 (preservatives)
Gmelin Reference	1443
KEGG	D05961
MeSH	Sulfur+dioxide
PubChem CID	1119
RTECS number	WS4550000
UNII	0UZA3422Q4
UN number	1079, 2037
CompTox Dashboard (EPA)	DTXSID6029672
InChI[show]
SMILES[show]
Properties
Chemical formula	SO 2
Molar mass	64.066 g mol−1
Appearance	Colorless gas
Odor	Pungent; similar to a just-struck match
Density	2.6288 kg m−3
Melting point	−72 °C; −98 °F; 201 K
Boiling point	−10 °C (14 °F; 263 K)
Solubility in water	94 g/L forms sulfurous acid
Vapor pressure	237.2 kPa
Acidity (pKa)	1.81
Basicity (pKb)	12.19
Magnetic susceptibility (χ)	−18.2·10−6 cm3/mol
Viscosity	0.403 cP (at 0 °C)
Structure
Point group	C2v
Coordination geometry	Digonal
Molecular shape	Dihedral
Dipole moment	1.62 D
Thermochemistry
Std molar entropy (So298)	248.223 J K−1 mol−1
Std enthalpy of formation (ΔfH⦵298)	−296.81 kJ mol−1
Hazards
GHS pictograms
GHS signal word	Danger
GHS hazard statements	H314, H331
NFPA 704	0 3 0
Lethal dose or concentration (LD, LC):
LC50 (median concentration)	3000 ppm (mouse, 30 min) 2520 ppm (rat, 1 hr)
LCLo (lowest published)	993 ppm (rat, 20 min) 611 ppm (rat, 5 hr) 764 ppm (mouse, 20 min) 1000 ppm (human, 10 min) 3000 ppm (human, 5 min)
US health exposure limits (NIOSH):
PEL (Permissible)	TWA 5 ppm (13 mg/m3)
REL (Recommended)	TWA 2 ppm (5 mg/m3) ST 5 ppm (13 mg/m3)
IDLH (Immediate danger)	100 ppm
Related compounds
Related sulfur oxides	Sulfur monoxide Sulfur trioxide
Related compounds	Ozone Selenium dioxide Sulfurous acid Tellurium dioxide
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Sulfur dioxide (also sulphur dioxide in British English) is the chemical compound with the formula SO
₂. It is a toxic gas responsible for the smell of burnt matches. It is released naturally by volcanic activity and is produced as a by-product of the burning of fossil fuels contaminated with sulfur compounds and copper extraction.

Structure and bonding

SO₂ is a bent molecule with C_2v 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. 

Two resonance structures of sulfur dioxide

 

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

The blue auroral glows of Io's upper atmosphere are caused by volcanic sulfur dioxide.

 

It is found on Earth and exists in very small concentrations and in the atmosphere at about 1 ppm.

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, 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. 

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.

Production

Sulfur dioxide is primarily produced for sulfuric acid manufacture. In the United States in 1979, 23.6 million tonnes (26,014,547 US short tons) of sulfur dioxide were used in this way, compared with 150 thousand tonnes (165,347 US 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:

S + O₂ → SO₂, ΔH = −297 kJ/mol

To aid combustion, liquefied sulfur (140–150 °C, 284-302 °F) 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, (1832-2912 °F). 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 (mercury sulfide) 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 concentrated 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 + Na₂SO₃ + H₂O

Reactions

Industrial reactions

Treatment of basic solutions with sulfur dioxide affords sulfite salts (e.g. sodium sulfite):

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 (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.

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 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. 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. 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 the risk of cork taint, 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.

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. 

It was shown that endogenous sulfur dioxide plays a role in diminishing an experimental lung damage caused by oleic acid. Endogenous sulfur dioxide lowered lipid peroxidation, free radical formation, oxidative stress and inflammation during an experimental lung damage. Conversely, a successful lung damage caused a significant lowering of endogenous sulfur dioxide production, and an increase in lipid peroxidation, free radical formation, oxidative stress and inflammation. Moreover, blockade of an enzyme that produces endogenous SO₂ significantly increased the amount of lung tissue damage in the experiment. Conversely, adding acetylcysteine or glutathione to the rat diet increased the amount of endogenous SO₂ produced and decreased the lung damage, the free radical formation, oxidative stress, inflammation and apoptosis.

It is considered that endogenous sulfur dioxide plays a significant physiological role in regulating cardiac and blood vessel function, and aberrant or deficient sulfur dioxide metabolism can contribute to several different cardiovascular diseases, such as arterial hypertension, atherosclerosis, pulmonary arterial hypertension, stenocardia.

It was shown that in children with pulmonary arterial hypertension due to congenital heart diseases the level of homocysteine is higher and the level of endogenous sulfur dioxide is lower than in normal control children. Moreover, these biochemical parameters strongly correlated to the severity of pulmonary arterial hypertension. Authors considered homocysteine to be one of useful biochemical markers of disease severity and sulfur dioxide metabolism to be one of potential therapeutic targets in those patients.

Endogenous sulfur dioxide also has been shown to lower the proliferation rate of endothelial smooth muscle cells in blood vessels, via lowering the MAPK activity and activating adenylyl cyclase and protein kinase A. Smooth muscle cell proliferation is one of important mechanisms of hypertensive remodeling of blood vessels and their stenosis, so it is an important pathogenetic mechanism in arterial hypertension and atherosclerosis. 

Endogenous sulfur dioxide in low concentrations causes endothelium-dependent vasodilation. In higher concentrations it causes endothelium-independent vasodilation and has a negative inotropic effect on cardiac output function, thus effectively lowering blood pressure and myocardial oxygen consumption. The vasodilating and bronchodilating effects of sulfur dioxide are mediated via ATP-dependent calcium channels and L-type ("dihydropyridine") calcium channels. Endogenous sulfur dioxide is also a potent antiinflammatory, antioxidant and cytoprotective agent. It lowers blood pressure and slows hypertensive remodeling of blood vessels, especially thickening of their intima. It also regulates lipid metabolism.

Endogenous sulfur dioxide also diminishes myocardial damage, caused by isoproterenol adrenergic hyperstimulation, and strengthens the myocardial antioxidant defense reserve.

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:

Proposed use in climate engineering

Injections of sulfur dioxide in the stratosphere has been proposed in climate engineering. The cooling effect would be similar to what has been observed after the large explosive volcano eruption of Mount Pinatubo in 1991. However this form of geoengineering would have uncertain regional consequences on rainfall patterns, for example in monsoon regions.

As an air pollutant

A sulfur dioxide plume from the Halemaʻumaʻu vent, glows at night

 

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: 

A collection of estimates of past and future anthropogenic global sulphur dioxide emissions. The Cofala et al. estimates are for sensitivity studies on SO₂ emission policies, CLE: Current Legislation, MFR: Maximum Feasible Reductions. RCPs (Representative Concentration Pathways) are used in CMIP5 simulations for latest (2013–2014) IPCC 5th assessment report.

 

Year	SO2
1970	31,161,000 short tons (28.3 Mt)
1980	25,905,000 short tons (23.5 Mt)
1990	23,678,000 short tons (21.5 Mt)
1996	18,859,000 short tons (17.1 Mt)
1997	19,363,000 short tons (17.6 Mt)
1998	19,491,000 short tons (17.7 Mt)
1999	18,867,000 short tons (17.1 Mt)

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 (lime) 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 burning by using limestone as a bed material in fluidized bed combustion.

Sulfur can also be removed from fuels before burning, preventing formation of SO₂ when the fuel is burnt. 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.

Sulfur dioxide in the world on April 15th, 2017. Note that sulfur dioxide moves through the atmosphere with prevailing winds and thus local sulfur dioxide distributions vary day to day with weather patterns and seasonality.

 

Fuel additives such as calcium additives and magnesium carboxylate may be used in marine 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 (23.1 Mt). This amount represents a 27% increase since 2000, and is roughly comparable with U.S. emissions in 1980.

Safety

US Geological Survey volunteer tests for sulfur dioxide after the 2018 lower Puna eruption

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 0.25 parts per million (ppm). The OSHA PEL is currently set at 5 ppm (13  mg/m³) 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 (ppb). 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.