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Sulfate



Sulfate


General
Molecular formula SO42-
Molar mass 96.06 g·mol−1
CAS number 14808-79-8
Structure
Molecular shape Tetrahedral
Related compounds
Related compounds Sulfonate
Except where noted otherwise, data are given for
materials in their standard state (at 25 °C, 100 kPa)
Infobox disclaimer and references

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

Contents

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:[1]

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. 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.[1] 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.[2]

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.[3] 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.[4][5] 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 dz2 and dx2-y2).[6] 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.[7][1] The Pauling bonding representation for sulfate and other main group compounds with oxygen is a common way of representing the bonding in many textbooks.[7][1]

Uses

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

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.[8] 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,[9] 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, [10] to facilitate cloud formation at low humidities and longer cloud lifetime.[11] 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.[12]

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 ion
SO42− Sulfate
SO32− Sulfite
S2O82− Peroxodisulfate
S2O72− Pyrosulfate
S2O62− Dithionate
S2O52− Metabisulfite
S2O42− Dithionite
S2O32− Thiosulfate
S4O62− Tetrathionate

See also

  • Sulfonate

References

  1. ^ a b c d Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements, 2nd Edition, Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4. 
  2. ^ The Atom and the Molecule by Gilbert N. Lewis Journal of the American Chemical Society Volume 38, 1916, pages 762-786
  3. ^ The modern theory of valency Linus Pauling J. Chem. Soc., 1948, 1461 - 1467, doi:10.1039/JR9480001461
  4. ^ C. A. Coulson, Nature, 221, 1106 (1969)
  5. ^ K. A. R. Mitchell, Chem. Rev., 69, 157 (1969)
  6. ^ Cotton, F. Albert; Wilkinson, Geoffrey (1966). Advanced Inorganic Chemistry (2d Edn.). New York:Wiley.
  7. ^ a b Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999). Advanced Inorganic Chemistry (6th Edn.) New York:Wiley-Interscience. ISBN 0-471-19957-5.
  8. ^ Inorganic and Theoretical Chemistry F.Sherwood Taylor 6th Edition (1942) William Heinemann
  9. ^ Figure 3: The global mean radiative forcing of the climate system for the year 2000, relative to 1750. Climate Change 2001: Working Group I: The Scientific Basis. IPCC.
  10. ^ Pincus & Baker 1994
  11. ^ Albrecht 1989
  12. ^ Chemical Amplification (or dampening) of the Twomey Effect: Conditions derived from droplet activation theory. T.A. Rissman, A. Nenes, J.H. Seinfeld.
 
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Sulfate". A list of authors is available in Wikipedia.
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