Sulfate
In inorganic chemistry, a sulfate (IUPAC-recommended spelling; also sulphate in British English) is a salt of sulfuric acid.
The structure and bonding of the sulfate ion
Chemical properties
The sulfate ion is a polyatomicanion with the empirical formulaSO2−
4 and a molecular mass of 96.06 daltons (96.06 g/mol); it consists of a central sulfuratom surrounded by four equivalent oxygen atoms in a tetrahedral arrangement. The symmetry is very similar to that of methane, CH4. The sulfur atom is in the +6 oxidation state while the four oxygen atoms are each in the −2 state. The sulfate ion carries a negative two charge and is the conjugate base of the bisulfate (or hydrogen sulfate) ion, HSO−
4, which is the conjugate base of H2SO4, sulfuric acid. Organic sulfates, such as dimethyl sulfate, are covalent compounds and esters of sulfuric acid.
Hydrogen sulfate (bisulfate)
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, that is no double bonds and a formal charge of 2+ on the sulfur atom.[1]
Two models of the sulfate ion.
1 with polar covalent bonds only. 2 with an ionic bond.
Later, Linus Pauling used valence bond theory to propose that the most significant resonance canonicals had two π bonds involving d orbitals. His reasoning was that the charge on sulfur was thus reduced, in accordance with his principle of electroneutrality.[2] 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.[3][4]
Double bonds in the Pauling structure imply a molecular orbital formed from 3d orbitals on sulfur and 2p orbitals on oxygen. A widely accepted description involving pπ - dπ bonding was initially proposed by D.W.J. 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 is 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 shorter S-O bonds have a different explanation. 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.[5][7]
Six resonances
The apparent contradiction can be cleared if one realizes that the covalent double bonds in the Lewis structure in reality represent bonds that are strongly polarized by more than 90% towards the oxygen atom. On the other hand, in the structure with an ionic bond, the charge is localized as a lone pair on the oxygen.[6]
Preparation
Methods of preparing ionic sulfates include:[7]
Properties
Many examples of ionic sulfates are known, and many of these are highly soluble in water. Exceptions include calcium sulfate, strontium sulfate, lead (II) 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.[7] An example is the neutral metal complex PtSO4(P(C6H5)3)2 where the sulfate ion is acting as a bidentate ligand. The metal-oxygen bonds in sulfate complexes can have significant covalent character.
Uses
Sulfates are important in both the chemical industry and biological systems:
Geometric Arrangement of Soap Bubbles
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 oxyanions
See also
Sulfonate
References
The structure and bonding of the sulfate ion
Chemical properties
The sulfate ion is a polyatomicanion with the empirical formulaSO2−
4 and a molecular mass of 96.06 daltons (96.06 g/mol); it consists of a central sulfuratom surrounded by four equivalent oxygen atoms in a tetrahedral arrangement. The symmetry is very similar to that of methane, CH4. The sulfur atom is in the +6 oxidation state while the four oxygen atoms are each in the −2 state. The sulfate ion carries a negative two charge and is the conjugate base of the bisulfate (or hydrogen sulfate) ion, HSO−
4, which is the conjugate base of H2SO4, sulfuric acid. Organic sulfates, such as dimethyl sulfate, are covalent compounds and esters of sulfuric acid.
Hydrogen sulfate (bisulfate)
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, that is no double bonds and a formal charge of 2+ on the sulfur atom.[1]
Two models of the sulfate ion.
1 with polar covalent bonds only. 2 with an ionic bond.
Later, Linus Pauling used valence bond theory to propose that the most significant resonance canonicals had two π bonds involving d orbitals. His reasoning was that the charge on sulfur was thus reduced, in accordance with his principle of electroneutrality.[2] 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.[3][4]
Double bonds in the Pauling structure imply a molecular orbital formed from 3d orbitals on sulfur and 2p orbitals on oxygen. A widely accepted description involving pπ - dπ bonding was initially proposed by D.W.J. 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 is 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 shorter S-O bonds have a different explanation. 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.[5][7]
Six resonances
The apparent contradiction can be cleared if one realizes that the covalent double bonds in the Lewis structure in reality represent bonds that are strongly polarized by more than 90% towards the oxygen atom. On the other hand, in the structure with an ionic bond, the charge is localized as a lone pair on the oxygen.[6]
Preparation
Methods of preparing ionic sulfates include:[7]
- 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, lead (II) 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.[7] An example is the neutral metal complex PtSO4(P(C6H5)3)2 where the sulfate ion is acting as a bidentate ligand. The metal-oxygen bonds in sulfate complexes can have significant covalent character.
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 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. These are known as sulfate-reducing bacteria.
- 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.
Geometric Arrangement of Soap Bubbles
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 oxyanions
| Molecular formula | Name |
|---|---|
| SO2− 5 | Peroxomonosulfate |
| SO2− 4 | Sulfate |
| SO2− 3 | Sulfite |
| S2O2− 8 | Peroxodisulfate |
| S2O2− 7 | Pyrosulfate |
| S2O2− 6 | Dithionate |
| S2O2− 5 | Metabisulfite |
| S2O2− 4 | Dithionite |
| S2O2− 3 | Thiosulfate |
| S4O2− 6 | Tetrathionate |
See also
Sulfonate
References
- ^The Atom and the Molecule by Gilbert N. Lewis Journal of the American Chemical Society Volume 38, 1916, pages 762-785. See page 778: 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 modern theory of valency Linus Pauling J. Chem. Soc., 1948, 1461 - 1467, doi:10.1039/JR9480001461
- ^ C.A. Coulson, Nature, 221, 1106 (1969)
- ^ K.A.R. Mitchell, Chem. Rev., 69, 157 (1969)
- ^ abCotton, F. Albert; Wilkinson, Geoffrey (1966). Advanced Inorganic Chemistry (2d Edn.). New York:Wiley.
- ^ ab Thorsten Stefan and Rudolf Janoschek: How relevant are S=O and P=O Double Bonds for the Description of the Acid Molecules H2SO3, H2SO4, and H3PO4, respectively?. Journal of Molecular Modeling, Volume 6, Number 2 / February 2000, p.282–288. doi:10.1007/PL00010730
- ^ abcGreenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. ISBN 0080379419.
- ^Inorganic and Theoretical Chemistry F.Sherwood Taylor 6th Edition (1942) William Heinemann
- ^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.
- ^ Pincus & Baker 1994
- ^ Albrecht 1989
- ^Chemical Amplification (or dampening) of the Twomey Effect: Conditions derived from droplet activation theory. T.A. Rissman, A. Nenes, J.H. Seinfeld.
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