Sulfate

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

Contents 1 Chemical properties 2 Uses 2.1 Oxoanion sulfides 3 Environmental effects 3.1 Main effects on climate 4 Notes

[edit] 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 one 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 hydrogensulfate ion, HSO₄⁻, which is the conjugate base of H₂SO₄, sulfuric acid.

Sulfate compounds arise when cations combine with the anion SO₄²⁻. Often this combination results in an ionic compound, although sulfates can engage in covalent bonding with most elements. The metal complex PtSO₄P(C₆H₅)₃₂ is clearly covalent Pt-O bonding. Dialkylsulfates, such as dimethylsulfate are covalent, distillable species. Many sulfate salts are highly soluble in water. Exceptions include calcium sulfate, strontium sulfate, and barium sulfate, which are poorly soluble. The barium derivative is useful in the gravimetric analysis of sulfate: one adds a solution of, say, barium chloride to a solution containing sulfate ions. The appearance of a white precipitate, which is barium sulfate, indicates that sulfate anions are present.

[edit] 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 utilize 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.

[edit] Oxoanion sulfides

• SO₅²⁻ persulfate ion
• SO₄²⁻ sulfate ion
• SO₃²⁻ sulfite ion
• SO₂²⁻ hyposulfite ion
• S₂O₈²⁻ peroxydisulfate ion

[edit] 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.

[edit] Main effects on climate

The first (direct) effect is to scatter 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,^[1] 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 (Pincus & Baker 1994), to facilitate cloud formation at low humidities and longer cloud lifetime (Albrecht 1989). 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 ^[2].

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.

[edit] Notes