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The sulfate or sulphate ion is a Polyatomic ion with the empirical formula . Salts, acid derivatives, and 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.
Later, Linus Pauling used valence bond theory to propose that the most significant resonance canonicals had two involving d orbitals. His reasoning was that the charge on sulfur was thus reduced, in accordance with his principle of electroneutrality. The S−O bond length of 149 pm is shorter than the bond lengths in sulfuric acid of 157 pm for S−OH. The double bonding was taken by Pauling to account for the shortness of the S−O bond.
Pauling's use of d orbitals provoked a debate on the relative importance of 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 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 d z2 and d x2– y2). 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).
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, Pauling's representation for sulfate and other main group compounds with oxygen is still a common way of representing the bonding in many textbooks. The apparent contradiction can be clarified if one realizes that the covalent bond double bonds in the Lewis structure actually represent bonds that are strongly polarized by more than 90% towards the oxygen atom. On the other hand, in the structure with a dipolar bond, the charge is localized as a lone pair on the oxygen.
Some metal can be oxidized to give metal sulfates.
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 complex or the neutral metal complex where the sulfate ion is acting as a denticity ligand. The metal–oxygen bonds in sulfate complexes can have significant covalent character.
Uses and occurrence
Commercial applications
Sulfates are widely used industrially. Major compounds include:
Occurrence in nature
Sulfate-reducing bacteria, 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.
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 iron(II) sulfate heptahydrate, ; blue vitriol is copper(II) sulfate pentahydrate, and white vitriol is zinc sulfate heptahydrate, . Alum, a double sulfate of potassium and aluminium with the formula , figured in the development of the chemical industry.
Environmental effects
Sulfates occur as microscopic particles (Particulate) resulting from fossil fuel and biomass combustion. They increase the acidity of the atmosphere and form acid rain. The anaerobic sulfate-reducing bacteria Desulfovibrio desulfuricans and D. vulgaris can remove the black sulfate crust that often tarnishes buildings.
Main effects on climate
Reversal and accelerated warming
Hydrological cycle
Solar geoengineering
As the real world had shown the importance of sulfate aerosol concentrations to the global climate, research into the subject accelerated. Formation of the aerosols and their effects on the atmosphere can be studied in the lab, with methods like ion-chromatography and mass spectrometry Samples of actual particles can be recovered from the stratosphere using balloons or aircraft, and remote were also used for observation. This data is fed into the , as the necessity of accounting for aerosol cooling to truly understand the rate and evolution of warming had long been apparent, with the IPCC Second Assessment Report being the first to include an estimate of their impact on climate, and every major model able to simulate them by the time IPCC Fourth Assessment Report was published in 2007. Many scientists also see the other side of this research, which is learning how to cause the same effect artificially. While discussed around the 1990s, if not earlier, stratospheric aerosol injection as a solar geoengineering method is best associated with Paul Crutzen's detailed 2006 proposal. Deploying in the stratosphere ensures that the aerosols are at their most effective, and that the progress of clean air measures would not be reversed: more recent research estimated that even under the highest-emission scenario RCP 8.5, the addition of stratospheric sulfur required to avoid relative to now (and relative to the preindustrial) would be effectively offset by the future controls on tropospheric sulfate pollution, and the amount required would be even less for less drastic warming scenarios. This spurred a detailed look at its costs and benefits, but even with hundreds of studies into the subject completed by the early 2020s, some notable uncertainties remain.
Hydrogensulfate (bisulfate)
The hydrogensulfate ion (), also called the bisulfate ion, is the conjugate base of sulfuric acid (). Sulfuric acid is classified as a strong acid; in aqueous solutions it ionizes completely to form hydronium () and hydrogensulfate () ions. In other words, the sulfuric acid behaves as a Brønsted–Lowry acid and is deprotonation to form hydrogensulfate ion. Hydrogensulfate has a valency of 1. An example of a salt containing the ion is sodium bisulfate, . In dilute solutions the hydrogensulfate ions also dissociate, forming more hydronium ions and sulfate ions ().
Other sulfur oxyanions
+Sulfur oxyanions
! Molecular formula
! Name Peroxomonosulfate Sulfate Sulfite Peroxydisulfate Pyrosulfate Dithionate Metabisulfite Dithionite Thiosulfate Trithionate Tetrathionate
See also
Notes
External links
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