Sulfate - Misplaced Pages

(Redirected from Sulfate salt) Anion of sulfur with 4 oxygen atoms This article is about the inorganic ion. For sulfate esters (commonly used in shampoo and personal care products), see Organosulfate.

Sulfate

Names

IUPAC name Sulfate

Other names Tetraoxosulfate(VI)
Tetraoxidosulfate(VI)

Identifiers

CAS Number

14808-79-8

3D model (JSmol)

Interactive image

ChEBI

CHEBI:16189

ChemSpider

1085

ECHA InfoCard

100.108.048

EC Number

233-334-2

PubChem CID

1117

UNII

7IS9N8KPMG

CompTox Dashboard (EPA)

DTXSID3042425

InChI

InChI=1S/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)/p-2Key: QAOWNCQODCNURD-UHFFFAOYSA-L
InChI=1/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)/p-2Key: QAOWNCQODCNURD-NUQVWONBAM

SMILES

S(=O)(=O)()

Properties

SO2−4

96.06 g·mol

Except where otherwise noted, data are given for materials in their standard state (at 25 °C , 100 kPa). Infobox references

Chemical compound

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.

Spelling

Further information: American and British English spelling differences

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

Structure

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, HSO−4, 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

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.

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. 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.

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_z and d_x–y). 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 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.

Preparation

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

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

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. An example is the complex Co(en)₂(SO₄)]Br 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

Commercial applications

Sulfates are widely used industrially. Major compounds include:

Gypsum, the natural mineral form of hydrated calcium sulfate, is used to produce plaster. About 100 million tonnes per year are used by the construction industry.
Copper sulfate, a common algaecide, the more stable form (CuSO₄) is used for galvanic cells as electrolyte
Iron(II) sulfate, a common form of iron in mineral supplements for humans, animals, and soil for plants
Magnesium sulfate (commonly known as Epsom salts), used in therapeutic baths
Lead(II) sulfate, produced on both plates during the discharge of a lead–acid battery
Sodium laureth sulfate, or SLES, a common detergent in shampoo formulations
Polyhalite, K₂Ca₂Mg(SO₄)₄·2H₂O, used as fertiliser.

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, 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

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.

Main effects on climate

This section is an excerpt from Global dimming § History.

Subsequent research estimated an average reduction in sunlight striking the terrestrial surface of around 4–5% per decade over the late 1950s–1980s, and 2–3% per decade when 1990s were included. Notably, solar radiation at the top of the atmosphere did not vary by more than 0.1-0.3% in all that time, strongly suggesting that the reasons for the dimming were on Earth. Additionally, only visible light and infrared radiation were dimmed, rather than the ultraviolet part of the spectrum. Further, the dimming had occurred even when the skies were clear, and it was in fact stronger than during the cloudy days, proving that it was not caused by changes in cloud cover alone.

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.

Reversal and accelerated warming

This section is an excerpt from Global dimming § Reversal.

After 1990, the global dimming trend had clearly switched to global brightening. 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. 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. By 2010, this reduction in sulfate pollution led to estimated healthcare cost savings valued at $50 billion annually. Similar measures were taken in Europe, 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.

At the peak of global dimming, it was able to counteract the warming trend completely. By 1975, the continually increasing concentrations of greenhouse gases have overcome the masking effect and dominated ever since. Even then, regions with high concentrations of sulfate aerosols due to air pollution had initially experienced cooling, in contradiction to the overall warming trend. 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.

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.

Regardless of the current strength of aerosol cooling, all future climate change scenarios project decreases in particulates and this includes the scenarios where 1.5 °C (2.7 °F) and 2 °C (3.6 °F) targets are met: their specific emission reduction targets assume the need to make up for lower dimming. Since models estimate that the cooling caused by sulfates is largely equivalent to the warming caused by atmospheric methane (and since methane is a relatively short-lived greenhouse gas), it is believed that simultaneous reductions in both would effectively cancel each other out.

Yet, in the recent years, methane concentrations had been increasing at rates exceeding their previous period of peak growth in the 1980s, with wetland methane emissions driving much of the recent growth, while air pollution is getting cleaned up aggressively. These trends are some of the main reasons why 1.5 °C (2.7 °F) warming is now expected around 2030, as opposed to the mid-2010s estimates where it would not occur until 2040.

Hydrological cycle

This section is an excerpt from Global dimming § Relationship with water cycle.

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. 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."

Solar geoengineering

This section is an excerpt from Stratospheric aerosol injection § Pollution controls and the discovery of radiative effects. 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 satellites were also used for observation. This data is fed into the climate models, 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 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. 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)

Hydrogensulfate
Names

IUPAC name Hydrogensulfate
Other names Bisulfate
Identifiers
CAS Number	14996-02-2
3D model (JSmol)	Interactive image
ChEBI	CHEBI:45696
ChemSpider	55666
ECHA InfoCard	100.108.048
Gmelin Reference	2121
PubChem CID	61778
CompTox Dashboard (EPA)	DTXSID3042425
InChI InChI=1S/H2O4S/c1-5(2,3)4/h(H2,1,2,3,4)/p-1Key: QAOWNCQODCNURD-UHFFFAOYSA-M
SMILES O(=O)(=O)
Properties
Chemical formula	HSO−4
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 , 100 kPa). Infobox references

Chemical compound

The hydrogensulfate ion (HSO−4), also called the bisulfate ion, is the conjugate base of sulfuric acid (H₂SO₄). Sulfuric acid is classified as a strong acid; in aqueous solutions it ionizes completely to form hydronium (H₃O) and hydrogensulfate (HSO−4) 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 HSO−4 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

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

Notes

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 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.