Sulfate
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Names | |||
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IUPAC name
Sulfate
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Other names
Tetraoxosulfate(VI)
Tetraoxidosulfate(VI) | |||
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ChEBI | |||
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ECHA InfoCard
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100.108.048 | ||
EC Number |
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PubChem CID
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UNII | |||
CompTox Dashboard (EPA)
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Properties | |||
SO2−4 | |||
Molar mass | 96.06 g·mol−1 | ||
Conjugate acid
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Hydrogensulfate
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Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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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
Bonding
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,
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 dx2–y2).[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.
Preparation
Typically
- 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
Uses and occurrence
Commercial applications
Sulfates are widely used industrially. Major compounds include:
- Gypsum, the natural mineral form of hydrated calcium sulfate, is used to produce plaster. About 100 million tonnes per year are used by the construction industry.
- Copper sulfate, a common algaecide, the more stable form (CuSO4) is used for galvanic cells as electrolyte
- Iron(II) sulfate, a common form of iron in mineral supplements for humans, animals, and soil for plants
- Epsom salts), used in therapeutic baths
- Lead(II) sulfate, produced on both plates during the discharge of a lead–acid battery
- Sodium laureth sulfate, or SLES, a common detergent in shampoo formulations
- fertiliser.
Occurrence in nature
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.
Environmental effects
Sulfates occur as microscopic particles (
Main effects on climate
Reversal and accelerated warming
After 1990, the global dimming trend had clearly switched to global brightening.
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
Hydrological cycle
Solar geoengineering
Hydrogensulfate (bisulfate)
Names | |
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IUPAC name
Hydrogensulfate[62]
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Other names
Bisulfate
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3D model (
JSmol ) |
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ChEBI | |
ChemSpider | |
ECHA InfoCard
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100.108.048 |
2121 | |
PubChem CID
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CompTox Dashboard (EPA)
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Properties | |
HSO−4 | |
Molar mass | 97.071 g/mol |
Conjugate acid
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Sulfuric acid |
Conjugate base
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Sulfate |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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The hydrogensulfate ion (HSO−4), 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 (HSO−4) 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 HSO−4 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
Molecular formula | Name |
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SO2−5 | Peroxomonosulfate |
SO2−4 | Sulfate |
SO2−3 | Sulfite |
S2O2−8 | Peroxydisulfate |
S2O2−7 | Pyrosulfate |
S2O2−6 | Dithionate |
S2O2−5 | Metabisulfite
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S2O2−4 | Dithionite |
S2O2−3 | Thiosulfate |
S3O2−6 | Trithionate |
S4O2−6 | Tetrathionate |
See also
- Sulfonate
- Sulfation and desulfation of lead–acid batteries
- Sulfate-reducing microorganisms
Notes
- ^ 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.
- ^ 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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