The sulfate or sulphate ion is a polyatomic anion with the empirical formula SO2−4. Salts, acid derivatives, and peroxides 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.

Sulfate
The structure and bonding of the sulfate ion. The distance between the sulfur atom and an oxygen atom is 149 picometers.
Ball-and-stick model of the sulfate anion
Names
IUPAC name
Sulfate
Other names
Tetraoxosulfate(VI)
Tetraoxidosulfate(VI)
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.108.048 Edit this at Wikidata
EC Number
  • 233-334-2
UNII
  • InChI=1S/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)/p-2
    Key: QAOWNCQODCNURD-UHFFFAOYSA-L
  • InChI=1/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)/p-2
    Key: QAOWNCQODCNURD-NUQVWONBAM
  • S(=O)(=O)([O-])[O-]
Properties
SO2−4
Molar mass 96.06 g·mol−1
Conjugate acid Hydrogensulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Spelling edit

"Sulfate" is the spelling recommended by IUPAC, but "sulphate" was traditionally used in British English.

Structure edit

The sulfate anion consists of a central sulfur atom surrounded by four equivalent oxygen atoms in a tetrahedral arrangement. The symmetry of the isolated anion is the same as that of methane. The sulfur atom is in the +6 oxidation state while the four oxygen atoms are each in the −2 state. The sulfate ion carries an overall charge of −2 and it is the conjugate base of the bisulfate (or hydrogensulfate) ion, HSO4, which is in turn the conjugate base of H2SO4, sulfuric acid. Organic sulfate esters, such as dimethyl sulfate, are covalent compounds and esters of sulfuric acid. The tetrahedral molecular geometry of the sulfate ion is as predicted by VSEPR theory.

Bonding edit

 
Two models of the sulfate ion.
1 with polar covalent bonds only; 2 with an ionic bond
 
Six resonances

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 and -1 on each oxygen atom.[1][a]

Later, Linus Pauling used valence bond theory to propose that the most significant resonance canonicals had two pi bonds involving d orbitals. His reasoning was that the charge on sulfur was thus reduced, in accordance with his principle of electroneutrality.[2] 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 pi 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]

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 dz2 and dx2y2).[5] 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).[6] 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, 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] 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 a dipolar bond, the charge is localized as a lone pair on the oxygen.[6]

Preparation edit

Typically metal sulfates are prepared by treating metal oxides, metal carbonates, or the metal itself with sulfuric acid:[7]

Zn + H2SO4 → ZnSO4 + H2
Cu(OH)2 + H2SO4 → CuSO4 + 2 H2O
CdCO3 + H2SO4 → CdSO4 + H2O + CO2

Although written with simple anhydrous formulas, these conversions generally are conducted in the presence of water. Consequently the product sulfates are hydrated, corresponding to zinc sulfate ZnSO4·7H2O, copper(II) sulfate CuSO4·5H2O, and cadmium sulfate CdSO4·H2O.

Some metal sulfides can be oxidized to give metal sulfates.

Properties edit

There are numerous examples of ionic sulfates, many of which are highly soluble in water. Exceptions include calcium sulfate, strontium sulfate, lead(II) sulfate, barium sulfate, silver sulfate, and mercury sulfate, which are poorly soluble. Radium sulfate is the most insoluble sulfate known. The barium derivative is useful in the gravimetric analysis of sulfate: if one adds a solution of most barium salts, for instance barium chloride, to a solution containing sulfate ions, barium sulfate will precipitate out of solution as a whitish powder. This is a common laboratory test to determine if 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 complex Co(en)2(SO4)]+Br[7] or the neutral metal complex PtSO4(PPh3)2] where the sulfate ion is acting as a bidentate ligand. The metal–oxygen bonds in sulfate complexes can have significant covalent character.

Uses and occurrence edit

Commercial applications edit

 
Knapsack sprayer used to apply sulfate to vegetables. Valencian Museum of Ethnology.

Sulfates are widely used industrially. Major compounds include:

Occurrence in nature edit

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 edit

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

Environmental effects edit

Sulfates occur as microscopic particles (aerosols) 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.[9]

Main effects on climate edit

 
This figure shows the level of agreement between a climate model driven by five factors and the historical temperature record. The negative component identified as "sulfate" is associated with the aerosol emissions blamed for global dimming.
 
The observed trends of global dimming and brightening in four major geopolitical regions. The dimming was greater on the average cloud-free days (red line) than on the average of all days (purple line), strongly suggesting that sulfate aerosols were the cause.[10]
In the 1990s, experiments comparing the atmosphere over the northern and southern islands of the Maldives, showed that the effect of macroscopic pollutants in the atmosphere at that time (blown south from India) caused about a 10% reduction in sunlight reaching the surface in the area under the Asian brown cloud – a much greater reduction than expected from the presence of the particles themselves.[11]
 
Sulfur dioxide in the world on April 15, 2017. Note that sulfur dioxide moves through the atmosphere with prevailing winds and thus local sulfur dioxide distributions vary day to day with weather patterns and seasonality.
Global dimming had been widely attributed to the increased presence of aerosol particles in Earth's atmosphere, predominantly those of sulfates.[12] While natural dust is also an aerosol with some impacts on climate, and volcanic eruptions considerably increase sulfate concentrations in the short term, these effects have been dwarfed by increases in sulfate emissions since the start of the Industrial Revolution.[13] According to the IPCC First Assessment Report, the global human-caused emissions of sulfur into the atmosphere were less than 3 million tons per year in 1860, yet they increased to 15 million tons in 1900, 40 million tons in 1940 and about 80 millions in 1980. This meant that the human-caused emissions became "at least as large" as all natural emissions of sulfur-containing compounds: the largest natural source, emissions of dimethyl sulfide from the ocean, was estimated at 40 million tons per year, while volcano emissions were estimated at 10 million tons. Moreover, that was the average figure: according to the report, "in the industrialized regions of Europe and North America, anthropogenic emissions dominate over natural emissions by about a factor of ten or even more".[14]

Reversal and accelerated warming edit

 
Sun-blocking aerosols around the world steadily declined (red line) since the 1991 eruption of Mount Pinatubo, according to satellite estimates.

After 1990, the global dimming trend had clearly switched to global brightening.[15][16][17][18][13] This followed measures taken to combat air pollution by the developed nations, typically through flue-gas desulfurization installations at thermal power plants, such as wet scrubbers or fluidized bed combustion.[19][20] In the United States, sulfate aerosols have declined significantly since 1970 with the passage of the Clean Air Act, which was strengthened in 1977 and 1990. According to the EPA, from 1970 to 2005, total emissions of the six principal air pollutants, including sulfates, dropped by 53% in the US.[21] By 2010, this reduction in sulfate pollution led to estimated healthcare cost savings valued at $50 billion annually.[22] Similar measures were taken in Europe,[21] such as the 1985 Helsinki Protocol on the Reduction of Sulfur Emissions under the Convention on Long-Range Transboundary Air Pollution, and with similar improvements.[23]

 
Satellite photo showing a thick pall of smoke and haze from forest fires in Eastern China. Such smoke is full of black carbon, which contributes to dimming trends but has an overall warming effect.
At the peak of global dimming, it was able to counteract the warming trend completely, but by 1975, the continually increasing concentrations of greenhouse gases have overcome the masking effect and dominated ever since.[21] Even then, regions with high concentrations of sulfate aerosols due to air pollution had initially experienced cooling, in contradiction to the overall warming trend.[24] The eastern United States was a prominent example: the temperatures there declined by 0.7 °C (1.3 °F) between 1970 and 1980, and by up to 1 °C (1.8 °F) in the Arkansas and Missouri. As the sulfate pollution was reduced, the central and eastern United States had experienced warming of 0.3 °C (0.54 °F) between 1980 and 2010,[25] even as sulfate particles still accounted for around 25% of all particulates.[22] By 2021, the northeastern coast of the United States was instead one of the fastest-warming regions of North America, as the slowdown of the Atlantic Meridional Overturning Circulation increased temperatures in that part of the North Atlantic Ocean.[26][27]

Since changes in aerosol concentrations already have an impact on the global climate, they would necessarily influence future projections as well. In fact, it is impossible to fully estimate the warming impact of all greenhouse gases without accounting for the counteracting cooling from aerosols. Climate models started to account for the effects of sulfate aerosols around the IPCC Second Assessment Report; when the IPCC Fourth Assessment Report was published in 2007, every climate model had integrated sulfates, but only 5 were able to account for less impactful particulates like black carbon.[28] By 2021, CMIP6 models estimated total aerosol cooling in the range from 0.1 °C (0.18 °F) to 0.7 °C (1.3 °F);[29] The IPCC Sixth Assessment Report selected the best estimate of a 0.5 °C (0.90 °F) cooling provided by sulfate aerosols, while black carbon amounts to about 0.1 °C (0.18 °F) of warming.[30] While these values are based on combining model estimates with observational constraints, including those on ocean heat content,[31] the matter is not yet fully settled. The difference between model estimates mainly stems from disagreements over the indirect effects of aerosols on clouds.[32][33]

It has also been suggested that aerosols are not given sufficient attention in regional risk assessments, in spite of being more influential on a regional scale than globally.[34] For instance, a climate change scenario with high greenhouse gas emissions but strong reductions in air pollution would see 0.2 °C (0.36 °F) more global warming by 2050 than the same scenario with little improvement in air quality, but regionally, the difference would add 5 more tropical nights per year in northern China and substantially increase precipitation in northern China and northern India.[35] Likewise, a paper comparing current level of clean air policies with a hypothetical maximum technically feasible action under otherwise the same climate change scenario found that the latter would increase the risk of temperature extremes by 30–50% in China and in Europe.[36] Unfortunately, because historical records of aerosols are sparser in some regions than in others, accurate regional projections of aerosol impacts are difficult. Even the latest CMIP6 climate models can only accurately represent aerosol trends over Europe,[10] but struggle with representing North America and Asia, meaning that their near-future projections of regional impacts are likely to contain errors as well.[37][10][38]

Hydrological cycle edit

 
Sulfate aerosols have decreased precipitation over most of Asia (red), but increased it over some parts of Central Asia (blue).[39]
On regional and global scale, air pollution can affect the water cycle, in a manner similar to some natural processes. One example is the impact of Sahara dust on hurricane formation: air laden with sand and mineral particles moves over the Atlantic Ocean, where they block some of the sunlight from reaching the water surface, slightly cooling it and dampening the development of hurricanes.[40] Likewise, it has been suggested since the early 2000s that since aerosols decrease solar radiation over the ocean and hence reduce evaporation from it, they would be "spinning down the hydrological cycle of the planet."[41][42] In 2011, it was found that anthropogenic aerosols had been the predominant factor behind 20th century changes in rainfall over the Atlantic Ocean sector,[43] when the entire tropical rain belt shifted southwards between 1950 and 1985, with a limited northwards shift afterwards.[44] Future reductions in aerosol emissions are expected to result in a more rapid northwards shift, with limited impact in the Atlantic but a substantially greater impact in the Pacific.[45]

Solar geoengineering edit

 
Proposed tethered balloon to inject aerosols into the stratosphere.
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[46] Samples of actual particles can be recovered from the stratosphere using balloons or aircraft, [47] and remote satellites were also used for observation.[48] This data is fed into the climate models,[49] 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.[50] Many scientists also see the other side of this research, which is learning how to cause the same effect artificially.[51] While discussed around the 1990s, if not earlier,[52] stratospheric aerosol injection as a solar geoengineering method is best associated with Paul Crutzen's detailed 2006 proposal.[53] 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 4 °C (7.2 °F) relative to now (and 5 °C (9.0 °F) 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.[54] This spurred a detailed look at its costs and benefits,[55] but even with hundreds of studies into the subject completed by the early 2020s, some notable uncertainties remain.[56]

Hydrogensulfate (bisulfate) edit

Hydrogensulfate
 
Names
IUPAC name
Hydrogensulfate[57]
Other names
Bisulfate
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.108.048  
2121
  • InChI=1S/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)/p-1
    Key: QAOWNCQODCNURD-UHFFFAOYSA-M
  • O[S](=O)(=O)[O-]
Properties
HSO4
Molar mass 97.071 g/mol
Conjugate acid Sulfuric acid
Conjugate base Sulfate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

The hydrogensulfate ion (HSO4), also called the bisulfate ion, is the conjugate base of sulfuric acid (H2SO4).[58][b] Sulfuric acid is classified as a strong acid; in aqueous solutions it ionizes completely to form hydronium (H3O+) and hydrogensulfate (HSO4) ions. In other words, the sulfuric acid behaves as a Brønsted–Lowry acid and is deprotonated to form hydrogensulfate ion. Hydrogensulfate has a valency of 1. An example of a salt containing the HSO4 ion is sodium bisulfate, NaHSO4. In dilute solutions the hydrogensulfate ions also dissociate, forming more hydronium ions and sulfate ions (SO2−4).

Other sulfur oxyanions edit

Sulfur oxyanions
Molecular formula Name
SO2−5 Peroxomonosulfate
SO2−4 Sulfate
SO2−3 Sulfite
S2O2−8 Peroxydisulfate
S2O2−7 Pyrosulfate
S2O2−6 Dithionate
S2O2−5 Metabisulfite
S2O2−4 Dithionite
S2O2−3 Thiosulfate
S3O2−6 Trithionate
S4O2−6 Tetrathionate

See also edit

Notes edit

  1. ^ 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.
  2. ^ The prefix "bi" in "bisulfate" comes from an outdated naming system and is based on the observation that there is twice as much sulfate (SO2−4) in sodium bisulfate (NaHSO4) and other bisulfates as in sodium sulfate (Na2SO4) and other sulfates. See also bicarbonate.

References edit

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