Sulfate

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Sulfate
The structure and bonding of the sulfate ion. The distance between the sulfur atom and an oxygen atom is 149 picometers.
Ball-and-stick model of the sulfate anion
Names
IUPAC name
Sulfate
Other names
Tetraoxosulfate(VI)
Tetraoxidosulfate(VI)
Identifiers
3D model (
JSmol
)
ChEBI
ChemSpider
ECHA InfoCard
100.108.048 Edit this at Wikidata
EC Number
  • 233-334-2
UNII
  • InChI=1S/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)/p-2
    Key: QAOWNCQODCNURD-UHFFFAOYSA-L
  • InChI=1/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)/p-2
    Key: QAOWNCQODCNURD-NUQVWONBAM
  • S(=O)(=O)([O-])[O-]
Properties
SO2−4
Molar mass 96.06 g·mol−1
Conjugate acid
Hydrogensulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

The sulfate or sulphate ion is a polyatomic anion with the empirical formula SO2−4. Salts, acid derivatives, and peroxides of sulfate are widely used in industry. Sulfates occur widely in everyday life. Sulfates are salts of sulfuric acid and many are prepared from that acid.

Spelling

"Sulfate" is the spelling recommended by IUPAC, but "sulphate" was traditionally used in British English.

Structure

The sulfate anion consists of a central

sulfate esters, such as dimethyl sulfate, are covalent compounds and esters of sulfuric acid. The tetrahedral molecular geometry of the sulfate ion is as predicted by VSEPR theory
.

Bonding

Six resonances

The first description of the bonding in modern terms was by Gilbert Lewis in his groundbreaking paper of 1916 where he described the bonding in terms of electron octets around each atom, that is no double bonds and a formal charge of +2 on the sulfur atom and -1 on each oxygen atom.[1][a]

Later,

electrostatic attraction) in causing the shortening of the S−O bond. The outcome was a broad consensus that d orbitals play a role, but are not as significant as Pauling had believed.[3][4]

A widely accepted description involving pπ – dπ bonding was initially proposed by Durward William John Cruickshank. In this model, fully occupied p orbitals on oxygen overlap with empty sulfur d orbitals (principally the dz2 and dx2y2).[5] However, in this description, despite there being some π character to the S−O bonds, the bond has significant ionic character. For sulfuric acid, computational analysis (with natural bond orbitals) confirms a clear positive charge on sulfur (theoretically +2.45) and a low 3d occupancy. Therefore, the representation with four single bonds is the optimal Lewis structure rather than the one with two double bonds (thus the Lewis model, not the Pauling model).[6] In this model, the structure obeys the octet rule and the charge distribution is in agreement with the electronegativity of the atoms. The discrepancy between the S−O bond length in the sulfate ion and the S−OH bond length in sulfuric acid is explained by donation of p-orbital electrons from the terminal S=O bonds in sulfuric acid into the antibonding S−OH orbitals, weakening them resulting in the longer bond length of the latter.

However, the bonding representation of Pauling for sulfate and other main group compounds with oxygen is still a common way of representing the bonding in many textbooks.

dipolar bond, the charge is localized as a lone pair on the oxygen.[6]

Preparation

Typically

metal sulfates are prepared by treating metal oxides, metal carbonates, or the metal itself with sulfuric acid:[7]

Zn + H2SO4 → ZnSO4 + H2
Cu(OH)2 + H2SO4 → CuSO4 + 2 H2O
CdCO3 + H2SO4 → CdSO4 + H2O + CO2

Although written with simple anhydrous formulas, these conversions generally are conducted in the presence of water. Consequently the product sulfates are hydrated, corresponding to zinc sulfate ZnSO4·7H2O, copper(II) sulfate CuSO4·5H2O, and cadmium sulfate CdSO4·H2O.

Some metal sulfides can be oxidized to give metal sulfates.

Properties

There are numerous examples of ionic sulfates, many of which are highly soluble in water. Exceptions include calcium sulfate, strontium sulfate, lead(II) sulfate, barium sulfate, silver sulfate, and mercury sulfate, which are poorly soluble. Radium sulfate is the most insoluble sulfate known. The barium derivative is useful in the gravimetric analysis of sulfate: if one adds a solution of most barium salts, for instance barium chloride, to a solution containing sulfate ions, barium sulfate will precipitate out of solution as a whitish powder. This is a common laboratory test to determine if sulfate anions are present.

The sulfate ion can act as a ligand attaching either by one oxygen (monodentate) or by two oxygens as either a

chelate or a bridge.[7] An example is the complex Co(en)2(SO4)]+Br[7] or the neutral metal complex PtSO4(PPh3)2] where the sulfate ion is acting as a bidentate
ligand. The metal–oxygen bonds in sulfate complexes can have significant covalent character.

Uses and occurrence

Commercial applications

Knapsack sprayer used to apply sulfate to vegetables. Valencian Museum of Ethnology.

Sulfates are widely used industrially. Major compounds include:

Occurrence in nature

Sulfate-reducing bacteria
, some anaerobic microorganisms, such as those living in sediment or near deep sea thermal vents, use the reduction of sulfates coupled with the oxidation of organic compounds or hydrogen as an energy source for chemosynthesis.

History

Some sulfates were known to alchemists. The vitriol salts, from the Latin vitreolum, glassy, were so-called because they were some of the first transparent crystals known.

white vitriol is zinc sulfate heptahydrate, ZnSO4·7H2O. Alum, a double sulfate of potassium and aluminium
with the formula K2Al2(SO4)4·24H2O, figured in the development of the chemical industry.

Environmental effects

Sulfates occur as microscopic particles (

D. vulgaris can remove the black sulfate crust that often tarnishes buildings.[9]

Main effects on climate

historical temperature record
. The negative component identified as "sulfate" is associated with the aerosol emissions blamed for global dimming.
The observed trends of global dimming and brightening in four major geopolitical regions. The dimming was greater on the average cloud-free days (red line) than on the average of all days (purple line), strongly suggesting that sulfate aerosols were the cause.[10]
Subsequent research estimated an average reduction in sunlight striking the terrestrial surface of around 4–5% per decade over late 1950s–1980s, and 2–3% per decade when 1990s were included.[11][12][13][14] Notably, solar radiation at the top of the atmosphere did not vary by more than 0.1-0.3% in all that time, strongly suggesting that the reasons for the dimming were on Earth.[15][16] Additionally, only visible light and infrared radiation were dimmed, rather than the ultraviolet part of the spectrum.[17] Further, the dimming had occurred even when the skies were clear, and it was in fact stronger than during the cloudy days, proving that it was not caused by changes in cloud cover alone.[18][16][10]
Sulfur dioxide in the world on April 15, 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.

Reversal and accelerated warming

aerosols around the world steadily declined (red line) since the 1991 eruption of Mount Pinatubo
, according to satellite estimates.

After 1990, the global dimming trend had clearly switched to global brightening.

EPA, from 1970 to 2005, total emissions of the six principal air pollutants, including sulfates, dropped by 53% in the US.[26] By 2010, this reduction in sulfate pollution led to estimated healthcare cost savings valued at $50 billion annually.[27] Similar measures were taken in Europe,[26] such as the 1985 Helsinki Protocol on the Reduction of Sulfur Emissions under the Convention on Long-Range Transboundary Air Pollution, and with similar improvements.[28]

forest fires in Eastern China. Such smoke is full of black carbon
, which contributes to dimming trends but has an overall warming effect.

Since changes in aerosol concentrations already have an impact on the global climate, they would necessarily influence future projections as well. In fact, it is impossible to fully estimate the warming impact of all

CMIP6 models estimated total aerosol cooling in the range from 0.1 °C (0.18 °F) to 0.7 °C (1.3 °F);[34] The IPCC Sixth Assessment Report selected the best estimate of a 0.5 °C (0.90 °F) cooling provided by sulfate aerosols, while black carbon amounts to about 0.1 °C (0.18 °F) of warming.[35] While these values are based on combining model estimates with observational constraints, including those on ocean heat content,[36] the matter is not yet fully settled. The difference between model estimates mainly stems from disagreements over the indirect effects of aerosols on clouds.[37][38]

It has also been suggested that aerosols are not given sufficient attention in regional risk assessments, in spite of being more influential on a regional scale than globally.
CMIP6 climate models can only accurately represent aerosol trends over Europe,[10] but struggle with representing North America and Asia, meaning that their near-future projections of regional impacts are likely to contain errors as well.[42][10][43]

Hydrological cycle

Sulfate aerosols have decreased precipitation over most of Asia (red), but increased it over some parts of Central Asia (blue).[44]
On regional and global scale, air pollution can affect the
solar radiation over the ocean and hence reduce evaporation from it, they would be "spinning down the hydrological cycle of the planet."[46][47] In 2011, it was found that anthropogenic aerosols had been the predominant factor behind 20th century changes in rainfall over the Atlantic Ocean sector,[48] when the entire tropical rain belt shifted southwards between 1950 and 1985, with a limited northwards shift afterwards.[49] Future reductions in aerosol emissions are expected to result in a more rapid northwards shift, with limited impact in the Atlantic but a substantially greater impact in the Pacific.[50]

Solar geoengineering

aerosols
into the stratosphere.
As the real world had shown the importance of sulfate aerosol concentrations to the global climate, research into the subject accelerated. Formation of the aerosols and their effects on the atmosphere can be studied in the lab, with methods like
Paul Crutzen's detailed 2006 proposal.[58] Deploying in the stratosphere ensures that the aerosols are at their most effective, and that the progress of clean air measures would not be reversed: more recent research estimated that even under the highest-emission scenario RCP 8.5, the addition of stratospheric sulfur required to avoid 4 °C (7.2 °F) relative to now (and 5 °C (9.0 °F) relative to the preindustrial) would be effectively offset by the future controls on tropospheric sulfate pollution, and the amount required would be even less for less drastic warming scenarios.[59] This spurred a detailed look at its costs and benefits,[60] but even with hundreds of studies into the subject completed by the early 2020s, some notable uncertainties remain.[61]

Hydrogensulfate (bisulfate)

Hydrogensulfate
Hydrogen sulfate (bisulfate)
Names
IUPAC name
Hydrogensulfate[62]
Other names
Bisulfate
Identifiers
3D model (
JSmol
)
ChEBI
ChemSpider
ECHA InfoCard
100.108.048 Edit this at Wikidata
2121
  • InChI=1S/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)/p-1
    Key: QAOWNCQODCNURD-UHFFFAOYSA-M
  • O[S](=O)(=O)[O-]
Properties
HSO4
Molar mass 97.071 g/mol
Conjugate acid
Sulfuric acid
Conjugate base
Sulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

The hydrogensulfate ion (HSO4), also called the bisulfate ion, is the conjugate base of sulfuric acid (H2SO4).[63][b] Sulfuric acid is classified as a strong acid; in aqueous solutions it ionizes completely to form hydronium (H3O+) and hydrogensulfate (HSO4) ions. In other words, the sulfuric acid behaves as a Brønsted–Lowry acid and is deprotonated to form hydrogensulfate ion. Hydrogensulfate has a valency of 1. An example of a salt containing the HSO4 ion is sodium bisulfate, NaHSO4. In dilute solutions the hydrogensulfate ions also dissociate, forming more hydronium ions and sulfate ions (SO2−4).

Other sulfur oxyanions

Sulfur oxyanions
Molecular formula Name
SO2−5 Peroxomonosulfate
SO2−4 Sulfate
SO2−3 Sulfite
S2O2−8 Peroxydisulfate
S2O2−7 Pyrosulfate
S2O2−6 Dithionate
S2O2−5
Metabisulfite
S2O2−4 Dithionite
S2O2−3 Thiosulfate
S3O2−6 Trithionate
S4O2−6 Tetrathionate

See also

Notes

  1. ^ Lewis assigned to sulfur a negative charge of two, starting from six own valence electrons and ending up with eight electrons shared with the oxygen atoms. In fact, sulfur donates two electrons to the oxygen atoms.
  2. ^ The prefix "bi" in "bisulfate" comes from an outdated naming system and is based on the observation that there is twice as much sulfate (SO2−4) in sodium bisulfate (NaHSO4) and other bisulfates as in sodium sulfate (Na2SO4) and other sulfates. See also bicarbonate.

References

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External links