sulfate, chemical compound containing the sulfate (SO₄) radical. Sulfates are salts or esters of sulfuric acid, H₂SO₄, formed by replacing one or both of the hydrogens with a metal (e.g., sodium) or a radical (e.g., ammonium or ethyl). Sulfates in which both hydrogens are replaced are called normal sulfates; sulfates in which only one hydrogen is replaced are called hydrogen sulfates, acid sulfates, or bisulfates. Most metal sulfates are readily soluble in water, but calcium and mercuric sulfates are only slightly soluble, while barium, lead, strontium, and mercurous sulfates are insoluble. In chemical analysis, the sulfate ion, SO₄^-2, is usually detected by adding barium chloride solution; the white barium sulfate precipitate that forms is insoluble in hydrochloric acid. Sulfates are widely distributed in nature. Barium sulfate occurs as barite; calcium sulfate is found as gypsum, alabaster, and selenite; Epsom salts is magnesium sulfate; sodium sulfate occurs as its decahydrate, Glauber's salt; and strontium sulfate occurs as celestite. Some sulfates were formerly known as vitriols; blue vitriol is cupric sulfate, green vitriol is ferrous sulfate, and white vitriol is zinc sulfate. Alums are double sulfates, containing two different metals and two sulfate radicals. Organic sulfates are esters. They can be formed by reacting an alcohol with cold sulfuric acid. They are also formed by the reaction of sulfuric acid with a double bond in an alkene; the product is called an alkyl hydrogen sulfate. An alkyl hydrogen sulfate can be broken down to an alcohol and sulfuric acid by heating it with water (hydrolysis); this reaction is often used to synthesize alcohols. Sulfates play a significant role both in the chemical industry and in biological systems. Sulfuric acid is used in lead storage batteries and in the manufacture of nitric acid; copper sulfate is a common algicide. Organisms found near deep-sea thermal vents use sulfates for energy in place of sunlight.

Any of numerous inorganic and organic chemical compounds related to sulfuric acid (H₂SO₄). One subgroup comprises salts containing the sulfate ion (SO₄²⁻) linked via ionic bonds with any of various cations. Another subgroup of sulfates, the esters, are organic compounds in which sulfuric acid's hydrogen atoms are replaced by organic groups (e.g., methyl, ethyl, phenyl); a carbon atom in the organic group bonds to an oxygen atom, whose second bond is to the sulfur atom. (In sulfonates, a carbon atom bonds directly to the sulfur atom.) Seealso bonding.

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or oil of vitriol

Dense, colourless, oily, corrosive liquid inorganic compound (H₂SO₄). A very strong acid, it forms ions of hydrogen or hydronium (H⁺ or H₃O⁺), hydrogen sulfate (HSO₄⁻), and sulfate (SO₄²⁻). It is also an oxidizing (see oxidation-reduction) and dehydrating agent and chars many organic materials. It is one of the most important industrial chemicals, used in various concentrations in manufacturing fertilizers, pigments, dyes, drugs, explosives, detergents, and inorganic salts and acids, in petroleum refining and metallurgical processes, and as the acid in lead-acid storage batteries. It is made industrially by dissolving sulfur trioxide (SO₃) in water, sometimes beyond the saturation point to make oleum (fuming sulfuric acid), used to make certain organic chemicals.

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In inorganic chemistry, a sulfate (IUPAC-recommended spelling; also sulphate in British English) is a salt of sulfuric acid.

Chemical properties

The sulfate ion is a polyatomic anion with the empirical formula SO₄²⁻ and a molecular mass of 96.06 daltons; it consists of a central sulfur atom surrounded by four equivalent oxygen atoms in a tetrahedral arrangement. The sulfate ion carries a negative two charge and is the conjugate base of the bisulfate (or hydrogen sulfate) ion, HSO₄⁻, which is the conjugate base of H₂SO₄, sulfuric acid. Organic sulfates, such as dimethyl sulfate, are covalent compounds and esters of sulfuric acid.

Preparation

Methods of preparing ionic sulfates include:

• dissolving a metal in sulfuric acid
• reacting sulfuric acid with a metal hydroxide or oxide
• oxidizing metal sulfides or sulfites

Properties

Many examples of ionic sulfates are known, and many of these are highly soluble in water. Exceptions include calcium sulfate, strontium sulfate, and barium sulfate, which are poorly soluble. Radium sulfate is the most insoluble sulfate known. The barium derivative is useful in the gravimetric analysis of sulfate: one adds a solution of, perhaps, barium chloride to a solution containing sulfate ions. The appearance of a white precipitate, which is barium sulfate, indicates that 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. An example is the neutral metal complex PtSO₄P(C₆H₅)₃₂ where the sulfate ion is acting as a bidentate ligand. The metal-oxygen bonds in sulfate complexes can have significant covalent character.

Structure and bonding

The S-O bond length of 149 pm is shorter than expected for a S-O single bond. For example, the bond lengths in sulfuric acid are 157 pm for S-OH. The tetrahedral geometry of the sulfate ion is as predicted by VSEPR theory.

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, i.e. no double bonds and a formal charge of 2+ on the sulfur atom.

Later, Linus Pauling used valence bond theory to propose that the most significant resonance canonicals had two π bonds (see above) involving d orbitals. His reasoning was that the charge on sulfur was thus reduced, in accordance with his principle of electroneutrality. The double bonding was taken by Pauling to account for the shortness of the S-O bond (149 pm).

Pauling's use of d orbitals provoked a debate on the relative importance of π bonding and bond polarity (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. A widely accepted description involves pπ - dπ bonding, initially proposed by D.W.J Cruickshank, where fully occupied p orbitals on oxygen overlap with empty sulfur d orbitals (principally the d_z² and d_x²-y²). In this description, while there is some π character to the S-O bonds, the bond has significant ionic character. This explanation is quoted in some current textbooks. The Pauling bonding representation for sulfate and other main group compounds with oxygen is a common way of representing the bonding in many textbooks.

Uses

Sulfates are important in both the chemical industry and biological systems:

• The lead-acid battery typically uses sulfuric acid.
• Some anaerobic microorganisms, such as those living near deep sea thermal vents use sulfates as electron acceptors.
• Copper sulfate is a common algaecide.
• Magnesium sulfate, commonly known as Epsom salts, is used in therapeutic baths.
• Gypsum, the natural mineral form of hydrated calcium sulfate, is used to produce plaster.
• The sulfate ion is used as counter ion for some cationic drugs.

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. Green vitriol is ferrous sulfate heptahydrate, FeSO₄·7H₂O; blue vitriol is copper sulfate pentahydrate, CuSO₄·5H₂O and white vitriol is zinc sulfate heptahydrate, ZnSO₄·7H₂O. Alum, a double sulfate with the formula K₂Al₂(SO₄)₄·24H₂O, figured in the development of the chemical industry.

Environmental effects

Sulfates occur as microscopic particles (aerosols) resulting from fossil fuel and biomass combustion. They increase the acidity of the atmosphere and form acid rain.

Main effects on climate

The main direct effect of sulfates on the climate involves the scattering of light, effectively increasing the Earth's albedo. This effect is moderately well understood and leads to a cooling from the negative radiative forcing of about 0.5 W/m² relative to pre-industrial values, partially offsetting the larger (about 2.4 W/m²) warming effect of greenhouse gases. The effect is strongly spatially non-uniform, being largest downstream of large industrial areas.

The first indirect effect is also known as the Twomey effect. Sulfate aerosols can act as cloud condensation nuclei and this leads to greater numbers of smaller droplets of water. Lots of smaller droplets can diffuse light more efficiently than just a few larger droplets.

The second indirect effect is the further knock-on effects of having more cloud condensation nuclei. It is proposed that these include the suppression of drizzle, increased cloud height, to facilitate cloud formation at low humidities and longer cloud lifetime. Sulfate may also result in changes in the particle size distribution, which can affect the clouds radiative properties in ways that are not fully understood. Chemical effects such as the dissolution of soluble gases and slightly soluble substances, surface tension depression by organic substances and accommodation coefficient changes are also included in the second indirect effect.

The indirect effects probably have a cooling effect, perhaps up to 2 W/m², although the uncertainty is very large. Sulfates are therefore implicated in global dimming, which may have acted to offset some of the effects of global warming.

Other sulfur oxoanions

Molecular formula	Name
SO52−	Peroxomonosulfate
SO42−	Sulfate
SO32−	Sulfite
S2O82−	Peroxodisulfate
S2O72−	Pyrosulfate
S2O62−	Dithionate
S2O52−	Metabisulfite
S2O42−	Dithionite
S2O32−	Thiosulfate
S4O62−	Tetrathionate

See also

• Sulfonate

References