The sulfate or sulphate ion is a
polyatomic anion with the
empirical formulaSO2−4. 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" 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.[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 dx2–y2).[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
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 sulfateZnSO4·7H2O,
copper(II) sulfateCuSO4·5H2O, and
cadmium sulfateCdSO4·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.[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
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.
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.[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.
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.[11][12][13][14] 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.[15][16] Additionally, only visible light and
infrared radiation were dimmed, rather than the
ultraviolet part of the spectrum.[17] 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.[18][16][10]
After 1990, the global dimming trend had clearly switched to global brightening.[19][20][21][22][23] 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.[24][25] 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.[26] By 2010, this reduction in sulfate pollution led to estimated healthcare cost savings valued at $50 billion annually.[27] Similar measures were taken in Europe,[26] 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.[28]
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.[26] Even then, regions with high concentrations of
sulfate aerosols due to air pollution had initially experienced cooling, in contradiction to the overall warming trend.[29] 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.[30]
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.[31][32]
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.[33] 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.[34][35] Yet, in the recent years, methane concentrations had been increasing at rates exceeding their previous period of peak growth in the 1980s,[36][37] with
wetland methane emissions driving much of the recent growth,[38][39] while air pollution is getting cleaned up aggressively.[40] 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.[31]
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
Saharadust 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.[42] 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."[43][44]
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]
The hydrogensulfate ion (HSO−4), 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 (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).
^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.
^
abStefan, Thorsten; Janoschek, Rudolf (Feb 2000). "How relevant are S=O and P=O Double Bonds for the Description of the Acid Molecules H2SO3, H2SO4, and H3PO4, respectively?". J. Mol. Modeling. 6 (2): 282–288.
doi:
10.1007/PL00010730.
S2CID96291857.
^Eddy, John A.; Gilliland, Ronald L.; Hoyt, Douglas V. (23 December 1982). "Changes in the solar constant and climatic effects". Nature. 300 (5894): 689–693.
Bibcode:
1982Natur.300..689E.
doi:
10.1038/300689a0.
S2CID4320853. Spacecraft measurements have established that the total radiative output of the Sun varies at the 0.1−0.3% level
^Lindeburg, Michael R. (2006). Mechanical Engineering Reference Manual for the PE Exam. Belmont, C.A.: Professional Publications, Inc. pp. 27–3.
ISBN978-1-59126-049-3.
^Risser, Mark D.; Collins, William D.; Wehner, Michael F.; O'Brien, Travis A.; Huang, Huanping; Ullrich, Paul A. (22 February 2024). "Anthropogenic aerosols mask increases in US rainfall by greenhouse gases". Nature Communications. 15.
doi:
10.1038/s41467-024-45504-8.
^Trisos, Christopher H.; Geden, Oliver; Seneviratne, Sonia I.; Sugiyama, Masahiro; van Aalst, Maarten; Bala, Govindasamy; Mach, Katharine J.; Ginzburg, Veronika; de Coninck, Heleen; Patt, Anthony (2021).
"Cross-Working Group Box SRM: Solar Radiation Modification"(PDF). Climate Change 2021: The Physical Science Basis. Contribution of Working Group I to the Sixth Assessment Report of the Intergovernmental Panel on Climate Change. 2021: 1238.
Bibcode:
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