Solar radiation management

From Wikipedia, the free encyclopedia

Proposed solar radiation management using a tethered balloon to inject sulfate aerosols into the stratosphere.

 

Solar radiation management (SRM) projects are a type of climate engineering which seek to reflect sunlight and thus reduce global warming. Proposed methods include increasing the planetary albedo, for example using stratospheric sulfate aerosols. Restorative methods have been proposed regarding the protection of natural heat reflectors like sea ice, snow and glaciers with engineering projects. Their principal advantages as an approach to climate engineering is the speed with which they can be deployed and become fully active, their potential low financial cost, and the reversibility of their direct climatic effects.

Solar radiation management projects could serve as a temporary response while levels of greenhouse gases can be brought under control by mitigation and greenhouse gas removal techniques. They would not reduce greenhouse gas concentrations in the atmosphere, and thus do not address problems such as ocean acidification caused by excess carbon dioxide (CO₂).

Purpose

Climate engineering projects have been proposed in order to reduce global warming. As early as 1974, Russian expert Mikhail Budyko suggested that if global warming became a problem, we could cool down the planet by burning sulfur in the stratosphere, which would create a haze. The annual cost of delivering a sufficient amount of sulfur to counteract expected greenhouse warming is estimated at 2 to 8 billion USD.

A preliminary study by Edward Teller and others in 1997 presented the pros and cons of various relatively "low-tech" proposals to mitigate global warming through scattering/reflecting sunlight away from the Earth via insertion of various materials in the upper stratosphere, low earth orbit, and L₁ locations.

By modifying the albedo of the Earth's surface, or by preventing sunlight reaching the Earth by using a solar shade, the sun's warming effect can be cancelled out—although the cancellation is imperfect, with regional discrepancies remaining. SRM or albedo modification, is considered to be a potential option for addressing climate change. As the National Academy of Sciences states in its 2015 report: "The two main options for responding to the risks of climate change involve mitigation—reducing and eventually eliminating human-caused emissions of CO₂ and other greenhouse gases (GHGs)—and adaptation—reducing the vulnerability of human and natural systems to changes in climate. A third potentially viable option, currently under development but not yet widely deployed, is carbon dioxide removal (CDR) from the atmosphere accompanied by reliable sequestration. A fourth, more speculative family of approaches called albedo modification seeks to offset climate warming by greenhouse gases by increasing the amount of sunlight reflected back to space." In this context, solar radiation management is widely viewed as a complement, not a substitute, to climate change mitigation and adaptation efforts. As The Royal Society concluded in its 2009 report: "Geoengineering methods are not a substitute for climate change mitigation, and should only be considered as part of a wider package of options for addressing climate change." Or put another way: "The safest and most predictable method of moderating climate change is to take early and effective action to reduce emissions of greenhouse gases. No geoengineering method can provide an easy or readily acceptable alternative solution to the problem of climate change. Geoengineering methods could, however, potentially be useful in future to augment continuing efforts to mitigate climate change by reducing emissions, and so should be subject to more detailed research and analysis."

By intentionally changing the Earth's albedo, or reflectivity, scientists propose that we could reflect more heat back out into space. We could also intercept sunlight before it reaches the Earth through a literal shade built in space. The effects are uncertain but it has been suggested that 2% albedo increase would roughly halve the effect of CO₂ doubling.

The National Academy of Sciences describes several of the potential benefits and risks of solar radiation management: "Modeling studies have shown that large amounts of cooling, equivalent in scale to the predicted warming due to doubling the CO₂ concentration in the atmosphere, can be produced by the introduction of tens of millions of tons of aerosols into the stratosphere. Preliminary modeling results suggest that albedo modification may be able to counter many of the damaging effects of high greenhouse gas concentrations on temperature and the hydrological cycle and reduce some impacts to sea ice. Models also strongly suggest that the benefits and risks will not be uniformly distributed around the globe."

The applicability of many techniques listed here has not been comprehensively tested. Even if the effects in computer simulation models or of small-scale interventions are known, there may be cumulative problems such as ozone depletion, which become apparent only from large-scale experiments.

Various small-scale experiments have been carried out on techniques such as cloud seeding, increasing the volume of stratospheric sulfate aerosols and implementing cool roof technology.

SRM has been suggested to control regional climate, but precise control over the geographical boundaries of the effect is not possible. 

SRM projects could, for example, be used as a temporary response while levels of greenhouse gases can be brought under control by greenhouse gas remediation techniques but would not reduce greenhouse gas concentrations in the atmosphere, and thus not address problems such as ocean acidification caused by excess carbon dioxide (CO₂).

Advantages

Solar radiation management has certain advantages relative to emissions cuts, adaptation, and carbon dioxide removal. Its effect of counteracting climate change would be experienced very rapidly, on the order of months after implementation, whereas the effects of emissions cuts and carbon dioxide removal are delayed because the climate change that they prevent is itself delayed. Some proposed solar radiation management techniques are expected to have very low direct financial costs of implementation, relative to the expected costs of both unabated climate change and aggressive mitigation. This creates a different problem structure. Whereas the provision of emissions reduction and carbon dioxide removal present collective action problems (because ensuring a lower atmospheric carbon dioxide concentration is a public good), a single countries or a handful of countries could implement solar radiation management. Finally, the direct climatic effects of solar radiation management are reversible on short timescales.

Limitations and risks

As well as the imperfect cancellation of the climatic effect of greenhouse gases, there are other significant problems with solar radiation management as a form of climate engineering. SRM is temporary in its effect, and thus and long-term restoration of the climate would rely on long-term SRM, unless carbon dioxide removal was subsequently used. However, short-term SRM programs are potentially beneficial.

Incomplete solution to CO₂ emissions

Solar radiation management does not remove greenhouse gases from the atmosphere and thus does not reduce other effects from these gases, such as ocean acidification. While not an argument against solar radiation management per se, this is an argument against reliance on climate engineering to the exclusion of greenhouse gas reduction.

Control and predictability

Most of the information on solar radiation management is from models and computer simulations. The actual results may differ from the predicted effect. The full effects of various solar radiation management proposals are not yet well understood. It may be difficult to predict the ultimate effects of projects, with models presently giving varying results. In the cases of systems which involve tipping points, effects may be irreversible. Furthermore, most modeling to date consider the effects of using solar radiation management to fully counteract the increase in global average surface temperature arising from a doubling or a quadrupling of the preindustrial carbon dioxide concentration. Under these assumptions, it overcompensates for the changes in precipitation from climate change. Solar radiation management is more likely to be optimized in a way that balances counteracting changes to temperature and precipitation, to compensate for some portion of climate change, and/or to slow down the rate of climate change.

Side effects

There may be unintended climatic consequences of solar radiation management, such as significant changes to the hydrological cycle that might not be predicted by the models used to plan them. Such effects may be cumulative or chaotic in nature. Ozone depletion is a risk of techniques involving sulfur delivery into the stratosphere. Not all side effects are negative, and an increase in agricultural productivity has been predicted by some studies due to the combination of more diffuse light and elevated carbon dioxide concentration.

Termination shock

If solar radiation management were masking a significant amount of warming and then were to abruptly stop, the climate would rapidly warm. This would cause a sudden rise in global temperatures towards levels which would have existed without the use of the climate engineering technique. The rapid rise in temperature may lead to more severe consequences than a gradual rise of the same magnitude.

Disagreement

Leaders of countries and other actors may disagree as to whether, how, and to what degree solar radiation would be used, which could exacerbate international tensions.

Weaponization

In 1976, 85 countries signed the U.N. Convention on the Prohibition of Military or Any Other Hostile Use of Environmental Modification Techniques. The Environmental Modification Convention generally prohibits weaponising climate engineering techniques. However, this does not eliminate the risk. If perfected to a degree of controllability and accuracy that is not considered possible at the moment, climate engineering techniques could theoretically be used by militaries to cause droughts or famines. Theoretically they could also be used simply to make battlefield conditions more favourable to one side or the other in a war.

Carnegie's Ken Caldeira said, "It will make it harder to achieve broad consensus on developing and governing these technologies if there is suspicion that gaining military advantage is an underlying motivation for its development..."

Effect on sunlight, sky and clouds

Managing solar radiation using aerosols or cloud cover would involve changing the ratio between direct and indirect solar radiation. This would affect plant life and solar energy. It is believed that there would be a significant effect on the appearance of the sky from stratospheric aerosol injection projects, notably a hazing of blue skies and a change in the appearance of sunsets. Aerosols affect the formation of clouds, especially cirrus clouds.

Proposed forms

Atmospheric

These projects seek to modify the atmosphere, either by enhancing naturally occurring stratospheric aerosols, or by using artificial techniques such as reflective balloons.

Stratospheric aerosols

Stratospheric Particle Injection for Climate Engineering

Injecting reflective aerosols into the stratosphere is the proposed solar radiation management method that has received the most sustained attention. This technique could give much more than 3.7 W/m² of globally averaged negative forcing, which is sufficient to entirely offset the warming caused by a doubling of CO₂, which is a common benchmark for assessing future climate scenarios. Sulfates are the most commonly proposed aerosols for climate engineering, since there is a good natural analogue with (and evidence from) volcanic eruptions. Explosive volcanic eruptions inject large amounts of sulfur dioxide gas into the stratosphere, which form sulfate aerosol and cool the planet. Alternative materials such as using photophoretic particles, titaniun dioxide, and diamond have been proposed. Delivery could be achieved using artillery, aircraft (such as the high-flying F15-C) or balloons. Broadly speaking, stratospheric aerosol injection is seen as a relatively more credible climate engineering technique, although one with potential major risks and challenges for its implementation. Risks include changes in precipitation and, in the case of sulfur, possible ozone depletion.

Marine cloud brightening

Various cloud reflectivity methods have been suggested, such as that proposed by John Latham and Stephen Salter, which works by spraying seawater in the atmosphere to increase the reflectivity of clouds. The extra condensation nuclei created by the spray will change the size distribution of the drops in existing clouds to make them whiter. The sprayers would use fleets of unmanned rotor ships known as Flettner vessels to spray mist created from seawater into the air to thicken clouds and thus reflect more radiation from the Earth. The whitening effect is created by using very small cloud condensation nuclei, which whiten the clouds due to the Twomey effect. 

This technique can give more than 3.7 W/m² of globally averaged negative forcing, which is sufficient to reverse the warming effect of a doubling of CO₂.

Ocean sulfur cycle enhancement

Enhancing the natural marine sulfur cycle by fertilizing a small portion with iron—typically considered to be a greenhouse gas remediation method—may also increase the reflection of sunlight. Such fertilization, especially in the Southern Ocean, would enhance dimethyl sulfide production and consequently cloud reflectivity. This could potentially be used as regional solar radiation management, to slow Antarctic ice from melting. Such techniques also tend to sequester carbon, but the enhancement of cloud albedo also appears to be a likely effect.

Terrestrial

Cool roof

The albedo of several types of roofs

Painting roof materials in white or pale colors to reflect solar radiation, known as 'cool roof' technology, is encouraged by legislation in some areas (notably California). This technique is limited in its ultimate effectiveness by the constrained surface area available for treatment. This technique can give between 0.01-0.19 W/m² of globally averaged negative forcing, depending on whether cities or all settlements are so treated. This is small relative to the 3.7 W/m² of positive forcing from a doubling of CO₂. Moreover, while in small cases it can be achieved at little or no cost by simply selecting different materials, it can be costly if implemented on a larger scale. A 2009 Royal Society report states that, "the overall cost of a 'white roof method' covering an area of 1% of the land surface (about 10¹² m²) would be about $300 billion/yr, making this one of the least effective and most expensive methods considered." However, it can reduce the need for air conditioning, which emits CO₂ and contributes to global warming.

Reflective sheeting

Adding reflective plastic sheets covering 67,000 square miles (170,000 km²) of desert every year between 2010 and 2070 to reflect the Sun's energy. may be able to give globally averaged 1.74 W/m² of negative forcing. Although insufficient to fully offset the 3.7 W/m² of positive forcing from a doubling of CO₂, this would still be a significant contribution thereto, and would offset the current level of warming (approx. 1.7 W/m²). However, the effect would be strongly regional, and would not be ideal for controlling Arctic shrinkage, which is one of the most significant problems resulting from global warming. Furthermore, desert albedo modification would be expensive, would compete with other land uses, and would have strongly negative ecological consequences. Finally, the total area required during 2010-70 is larger than all non-polar deserts combined.

Ocean changes

Oceanic foams have also been suggested, using microscopic bubbles suspended in the upper layers of the photic zone. A less costly proposal is to simply lengthen and brighten existing ship wakes.

Ice protection

Arctic sea ice formation could be increased by pumping deep cooler water to the surface. Sea ice (and terrestrial) ice can be thickened by increasing albedo with silica spheres. Glaciers flowing into the sea may be stabilized by blocking the flow of warm water to the glacier.

Forestry

Reforestation in tropical areas has a cooling effect. Deforestation of high-latitude and high-altitude forests exposes snow and this increases albedo.

Grassland management

Changes to grassland have been proposed to increase albedo. This technique can give 0.64 W/m² of globally averaged negative forcing, which is insufficient to offset the 3.7 W/m² of positive forcing from a doubling of CO₂, but could make a minor contribution.

High-albedo crop varieties

Selecting or genetically modifying commercial crops with high albedo has been suggested. This has the advantage of being relatively simple to implement, with farmers simply switching from one variety to another. Temperate areas may experience a 1 °C cooling as a result of this technique. This technique is an example of bio-geoengineering. This technique can give 0.44 W/m² of globally averaged negative forcing, which is insufficient to offset the 3.7 W/m² of positive forcing from a doubling of CO₂, but could make a minor contribution.

Space-based

Space-based climate engineering projects (space sunshades) are seen by many commentators and scientists as being very expensive and technically difficult, with the Royal Society suggesting that "the costs of setting in place such a space-based armada for the relatively short period that SRM geoengineering may be considered applicable (decades rather than centuries) would likely make it uncompetitive with other SRM approaches."

Space mirrors

Proposed by Roger Angel with the purpose to deflect a percentage of solar sunlight into space, using mirrors orbiting around the Earth.

Moon dust

Mining moon dust to create a shielding cloud was proposed by Curtis Struck at Iowa State University in Ames.

Dispersive solutions

The basic function of a space lens to mitigate global warming. In reality, a 1000 kilometre diameter lens is enough, much smaller than what is shown in the simplified image. In addition, as a Fresnel lens it would only be a few millimeters thick.

 

Several authors have proposed dispersing light before it reaches the Earth by putting a very large diffraction grating (thin wire mesh) or lens in space, perhaps at the L1 point between the Earth and the Sun. Using a Fresnel lens in this manner was proposed in 1989 by J. T. Early. Using a diffraction grating was proposed in 1997 by Edward Teller, Lowell Wood, and Roderick Hyde. In 2004, physicist and science fiction author Gregory Benford calculated that a concave rotating Fresnel lens 1000 kilometers across, yet only a few millimeters thick, floating in space at the L₁ point, would reduce the solar energy reaching the Earth by approximately 0.5% to 1%. He estimated that this would cost around US$10 billion up front, and another $10 billion in supportive cost during its lifespan. One issue with implementing such a solution is the need to counteract the effects of the solar wind moving such megastructures out of position.

Governance

Climate engineering poses several challenges in the context of governance because of issues of power and jurisdiction. Climate engineering as a climate change solution differs from other mitigation and adaptation strategies. Unlike a carbon trading system that would be focused on participation from multiple parties along with transparency, monitoring measures and compliance procedures; this is not necessarily required by climate engineering. Bengtsson (2006) argues that "the artificial release of sulphate aerosols is a commitment of at least several hundred years". Yet this is true only if a long-term deployment strategy is adopted. Under a short-term, temporary strategy, implementation would instead be limited to decades. Both cases, however, highlight the importance for a political framework that is sustainable enough to contain a multilateral commitment over such a long period and yet is flexible as the techniques innovate through time. There are many controversies surrounding this topic and hence, climate engineering has been made into a very political issue. Most discussions and debates are not about which climate engineering technique is better than the other, or which one is more economically and socially feasible. Discussions are broadly on who will have control over the deployment of climate engineering and under what governance regime the deployment can be monitored and supervised. This is especially important due to the regional variability of the effects of many climate engineering techniques, benefiting some countries while damaging others. The challenge posed by climate engineering is not how to get countries to do it. It is to address the fundamental question of who should decide whether and how climate engineering should be attempted – a problem of governance.

Solar radiation management raises a number of governance challenges. David Keith argues that the cost is within the realm of small countries, large corporations, or even very wealthy individuals. David Victor suggests that climate engineering is within the reach of a lone "Greenfinger," a wealthy individual who takes it upon him or herself to be the "self-appointed protector of the planet". However, it has been argued that a rogue state threatening solar radiation management may strengthen action on mitigation.

Legal and regulatory systems may face a significant challenge in effectively regulating solar radiation management in a manner that allows for an acceptable result for society. There are, however, significant incentives for states to cooperate in choosing a specific climate engineering policy, which make unilateral deployment a rather unlikely event.

Some researchers have suggested that building a global agreement on climate engineering deployment will be very difficult, and instead power blocs are likely to emerge.

Public attitudes

There have been a handful of studies into attitudes to and opinions of solar radiation management. These generally find low levels of awareness, uneasiness with the implementation of solar radiation management, cautious support of research, and a preference for greenhouse gas emissions reduction. As is often the case with public opinions regarding emerging issues, the responses are highly sensitive to the questions' particular wording and context. 

One cited objection to implementing a short-term temperature fix is that there might then be less incentive to reduce carbon dioxide emissions until it caused some other environmental catastrophe, such as a chemical change in ocean water that could be disastrous to ocean life.

Stratospheric sulfur aerosols

From Wikipedia, the free encyclopedia

Stratospheric sulfur aerosols are sulfur-rich particles which exist in the stratosphere region of the Earth's atmosphere. The layer of the atmosphere in which they exist is known as the Junge layer, or simply the stratospheric aerosol layer. These particles consist of a mixture of sulfuric acid and water. They are created naturally, such as by photochemical decomposition of sulfur-containing gases, e.g. carbonyl sulfide. When present in high levels, e.g. after a strong volcanic eruption such as Mount Pinatubo, they produce a cooling effect, by reflecting sunlight, and by modifying clouds as they fall out of the stratosphere. This cooling may persist for a few years before the particles fall out.

An aerosol is a suspension of fine solid particles or liquid droplets in a gas. The sulfate particles or sulfuric acid droplets in the atmosphere are about 0.1 to 1.0 micrometer (a millionth of a meter) in diameter.

Sulfur aerosols are common in the troposphere as a result of pollution with sulfur dioxide from burning coal, and from natural processes. Volcanos are a major source of particles in the stratosphere as the force of the volcanic eruption propels sulfur-containing gases into the stratosphere. The relative influence of volcanoes on the Junge layer varies considerably according to the number and size of eruptions in any given time period, and also of quantities of sulfur compounds released. Only stratovolcanoes containing primarily granitic rocks are responsible for these fluxes, as basaltic rock erupted in shield volcanoes doesn't result in plumes which reach the stratosphere.

Creating stratospheric sulfur aerosols deliberately is a proposed geoengineering technique which offers a possible solution to some of the problems caused by global warming. However, this will not be without side effects  and it has been suggested that the cure may be worse than the disease.

Pinatubo eruption cloud. This volcano released huge quantities of stratospheric sulfur aerosols and contributed greatly to understanding of the subject.

Origins

Volcanic "injection"

 

Natural sulfur aerosols are formed in vast quantities from the SO₂ ejected by volcanoes, which may be injected directly into the stratosphere during very large (Volcanic Explosivity Index, VEI, of 4 or greater) eruptions. A comprehensive analysis, dealing largely with tropospheric sulfur compounds in the atmosphere, is provided by Bates et al.

The IPCC AR4 says explosive volcanic events are episodic, but the stratospheric aerosols resulting from them yield substantial transitory perturbations to the radiative energy balance of the planet, with both shortwave and longwave effects sensitive to the microphysical characteristics of the aerosols.

During periods lacking volcanic activity (and thus direct injection of SO₂ into the stratosphere), oxidation of COS (carbonyl sulfide) dominates the production of stratospheric sulfur aerosol.

Chemistry

The chemistry of stratospheric sulfur aerosols varies significantly according to their source. Volcanic emissions vary significantly in composition, and have complex chemistry due to the presence of ash particulates and a wide variety of other elements in the plume.

The chemical reactions affecting both the formation and elimination of sulfur aerosols are not fully understood. It is difficult to estimate accurately, for example, whether the presence of ash and water vapour is important for aerosol formation from volcanic products, and whether high or low atmospheric concentrations of precursor chemicals (such as SO₂ and H₂S) are optimal for aerosol formation. This uncertainty makes it difficult to determine a viable approach for geoengineering uses of sulfur aerosol formation.

Scientific study

Stratospheric sulfates from volcanic emissions cause transient cooling; the purple line showing sustained cooling is from tropospheric sulfate

 

Understanding of these aerosols comes in large part from the study of volcanic eruptions, notably Mount Pinatubo in the Philippines, which erupted in 1991 when scientific techniques were sufficiently far advanced to study the effects carefully.

The formation of the aerosols and their effects on the atmosphere can also be studied in the lab. Samples of actual particles can be recovered from the stratosphere using balloons or aircraft.

Computer models can be used to understand the behavior of aerosol particles, and are particularly useful in modelling their effect on global climate.

Biological experiments in the lab, and field/ocean measurements can establish the formation mechanisms of biologically derived volatile sulfurous gases.

Effects

It has been established that emission of precursor gases for sulfur aerosols is the principal mechanism by which volcanoes cause episodic global cooling. The Intergovernmental Panel on Climate Change AR4 regards stratospheric sulfate aerosols as having a low level of scientific understanding. The aerosol particles form a whitish haze in the sky. This creates a global dimming effect, where less of the sun's radiation is able to reach the surface of the Earth. This leads to a global cooling effect. In essence, they act as the reverse of a greenhouse gas, which tends to allow visible light from the sun through, whilst blocking infrared light emitted from the Earth's surface and its atmosphere. The particles also radiate infrared energy directly, as they lose heat into space. 

Solar radiation reduction due to volcanic eruptions

 

All aerosols both absorb and scatter solar and terrestrial radiation. This is quantified in the Single Scattering Albedo (SSA), the ratio of scattering alone to scattering plus absorption (extinction) of radiation by a particle. The SSA tends to unity if scattering dominates, with relatively little absorption, and decreases as absorption increases, becoming zero for infinite absorption. For example, sea-salt aerosol has an SSA of 1, as a sea-salt particle only scatters, whereas soot has an SSA of 0.23, showing that it is a major atmospheric aerosol absorber.

Aerosols, natural and anthropogenic, can affect the climate by changing the way radiation is transmitted through the atmosphere. Direct observations of the effects of aerosols are quite limited so any attempt to estimate their global effect necessarily involves the use of computer models. The Intergovernmental Panel on Climate Change, IPCC, says: While the radiative forcing due to greenhouse gases may be determined to a reasonably high degree of accuracy... the uncertainties relating to aerosol radiative forcings remain large, and rely to a large extent on the estimates from global modelling studies that are difficult to verify at the present time. However, they are mostly talking about tropospheric aerosol. 

The aerosols have a role in the destruction of ozone due to surface chemistry effects. Destruction of ozone has in recent years created large holes in the ozone layer, initially over the Antarctic and then the Arctic. These holes in the ozone layer have the potential to expand to cover inhabited and vegetative regions of the planet, leading to catastrophic environmental damage. 

Ozone destruction occurs principally in polar regions, but the formation of ozone occurs principally in the tropics. Ozone is distributed around the planet by the Brewer-Dobson circulation. Therefore, the source and dispersal pattern of aerosols is critical in understanding their effect on the ozone layer. 

Turner was inspired by dramatic sunsets caused by volcanic aerosols

 

Aerosols scatter light, which affects the appearance of the sky and of sunsets. Changing the concentration of aerosols in the atmosphere can dramatically affect the appearance of sunsets. A change in sky appearance during 1816, "The Year Without A Summer" (attributed to the eruption of Mount Tambora), was the inspiration for the paintings of J. M. W. Turner. Further volcanic eruptions and geoengineering projects involving sulfur aerosols are likely to affect the appearance of sunsets significantly, and to create a haze in the sky. 

Aerosol particles are eventually deposited from the stratosphere onto land and ocean. Depending on the volume of particles descending, the effects may be significant to ecosystems, or may not be. Modelling of the quantities of aerosols used in likely geoengineering scenarios suggest that effects on terrestrial ecosystems from deposition is not likely to be significantly harmful.

Climate engineering

The ability of stratospheric sulfur aerosols to create this global dimming effect has made them a possible candidate for use in climate engineering projects to limit the effect and impact of climate change due to rising levels of greenhouse gases. Delivery of precursor gases such as H₂S and SO₂ by artillery, aircraft and balloons has been proposed.

Understanding of this proposed technique is partly based on the fact that it is the adaptation of an existing atmospheric process. The technique is therefore potentially better understood than are comparable (but purely speculative) climate engineering proposals. It is also partly based on the speed of action of any such solution deployed, in contrast to carbon sequestration projects such as carbon dioxide air capture which would take longer to work. However, gaps in understanding of these processes exist, for example the effect on stratospheric climate and on rainfall patterns, and further research is needed.

Sulfur

From Wikipedia, the free encyclopedia

Sulfur,  ₁₆S

General properties
Alternative name	sulphur (British spelling)
Appearance	lemon yellow sintered microcrystals
Standard atomic weight (Ar, standard)	[32.059, 32.076] conventional: 32.06
Sulfur in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson O ↑ S ↓ Se phosphorus ← sulfur → chlorine
Atomic number (Z)	16
Group	group 16 (chalcogens)
Period	period 3
Block	p-block
Element category	reactive nonmetal
Electron configuration	[Ne] 3s2 3p4
Electrons per shell	2, 8, 6
Physical properties
Phase at STP	solid
Melting point	388.36 K ​(115.21 °C, ​239.38 °F)
Boiling point	717.8 K ​(444.6 °C, ​832.3 °F)
Density (near r.t.)	alpha: 2.07 g/cm3 beta: 1.96 g/cm3 gamma: 1.92 g/cm3
when liquid (at m.p.)	1.819 g/cm3
Critical point	1314 K, 20.7 MPa
Heat of fusion	mono: 1.727 kJ/mol
Heat of vaporization	mono: 45 kJ/mol
Molar heat capacity	22.75 J/(mol·K)
Vapor pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 375 408 449 508 591 717
Atomic properties
Oxidation states	−2, −1, +1, +2, +3, +4, +5, +6 (a strongly acidic oxide)
Electronegativity	Pauling scale: 2.58
Ionization energies	1st: 999.6 kJ/mol 2nd: 2252 kJ/mol 3rd: 3357 kJ/mol
Covalent radius	105±3 pm
Van der Waals radius	180 pm
Spectral lines of sulfur
Other properties
Natural occurrence	primordial
Crystal structure	​orthorhombic
Thermal conductivity	0.205 W/(m·K) (amorphous)
Electrical resistivity	2×1015  Ω·m (at 20 °C) (amorphous)
Magnetic ordering	diamagnetic
Magnetic susceptibility	(α) −15.5·10−6 cm3/mol (298 K)
Bulk modulus	7.7 GPa
Mohs hardness	2.0
CAS Number	7704-34-9
History
Discovery	Chinese (before 2000 BCE)
Recognized as an element by	Antoine Lavoisier (1777)
Main isotopes of sulfur
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct 32S 94.99% stable 33S 0.75% stable 34S 4.25% stable 35S trace 87.32 d β− 35Cl 36S 0.01% stable

Sulfur or sulphur is a chemical element with symbol S and atomic number 16. It is abundant, multivalent, and nonmetallic. Under normal conditions, sulfur atoms form cyclic octatomic molecules with a chemical formula S₈. Elemental sulfur is a bright yellow, crystalline solid at room temperature.
Sulfur is the tenth most common element by mass in the universe, and the fifth most common on Earth. Though sometimes found in pure, native form, sulfur on Earth usually occurs as sulfide and sulfate minerals. Being abundant in native form, sulfur was known in ancient times, being mentioned for its uses in ancient India, ancient Greece, China, and Egypt. In the Bible, sulfur is called brimstone. Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum. The greatest commercial use of the element is the production of sulfuric acid for sulfate and phosphate fertilizers, and other chemical processes. The element sulfur is used in matches, insecticides, and fungicides. Many sulfur compounds are odoriferous, and the smells of odorized natural gas, skunk scent, grapefruit, and garlic are due to organosulfur compounds. Hydrogen sulfide gives the characteristic odor to rotting eggs and other biological processes.

Sulfur is an essential element for all life, but almost always in the form of organosulfur compounds or metal sulfides. Three amino acids (cysteine, cystine, and methionine) and two vitamins (biotin and thiamine) are organosulfur compounds. Many cofactors also contain sulfur including glutathione and thioredoxin and iron–sulfur proteins. Disulfides, S–S bonds, confer mechanical strength and insolubility of the protein keratin, found in outer skin, hair, and feathers. Sulfur is one of the core chemical elements needed for biochemical functioning and is an elemental macronutrient for all living organisms.

Characteristics

When burned, sulfur melts to a blood-red liquid and emits a blue flame that is best observed in the dark.

Physical properties

Sulfur forms polyatomic molecules with different chemical formulas, the best-known allotrope being octasulfur, cyclo-S₈. The point group of cyclo-S₈ is D_4d and its dipole moment is 0 D. Octasulfur is a soft, bright-yellow solid that is odorless, but impure samples have an odor similar to that of matches. It melts at 115.21 °C (239.38 °F), boils at 444.6 °C (832.3 °F) and sublimes easily. At 95.2 °C (203.4 °F), below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-polymorph. The structure of the S₈ ring is virtually unchanged by this phase change, which affects the intermolecular interactions. Between its melting and boiling temperatures, octasulfur changes its allotrope again, turning from β-octasulfur to γ-sulfur, again accompanied by a lower density but increased viscosity due to the formation of polymers. At higher temperatures, the viscosity decreases as depolymerization occurs. Molten sulfur assumes a dark red color above 200 °C (392 °F). The density of sulfur is about 2 g/cm³, depending on the allotrope; all of the stable allotropes are excellent electrical insulators.

Chemical properties

Sulfur burns with a blue flame with formation of sulfur dioxide, which has a suffocating and irritating odor. Sulfur is insoluble in water but soluble in carbon disulfide and, to a lesser extent, in other nonpolar organic solvents, such as benzene and toluene. The first and second ionization energies of sulfur are 999.6 and 2252 kJ/mol, respectively. Despite such figures, the +2 oxidation state is rare, with +4 and +6 being more common. The fourth and sixth ionization energies are 4556 and 8495.8 kJ/mol, the magnitude of the figures caused by electron transfer between orbitals; these states are only stable with strong oxidants such as fluorine, oxygen, and chlorine. Sulfur reacts with nearly all other elements with the exception of the noble gases, even with the notoriously unreactive metal iridium (yielding iridium disulfide). Some of those reactions need elevated temperatures.

Allotropes

The structure of the cyclooctasulfur molecule, S₈

Sulfur forms over 30 solid allotropes, more than any other element. Besides S₈, several other rings are known. Removing one atom from the crown gives S₇, which is more of a deep yellow than the S₈. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S₈, but with S₇ and small amounts of S₆. Larger rings have been prepared, including S₁₂ and S₁₈.

Amorphous or "plastic" sulfur is produced by rapid cooling of molten sulfur—for example, by pouring it into cold water. X-ray crystallography studies show that the amorphous form may have a helical structure with eight atoms per turn. The long coiled polymeric molecules make the brownish substance elastic, and in bulk this form has the feel of crude rubber. This form is metastable at room temperature and gradually reverts to crystalline molecular allotrope, which is no longer elastic. This process happens within a matter of hours to days, but can be rapidly catalyzed.

Isotopes

Sulfur has 25 known isotopes, four of which are stable: ³²S (94.99%±0.26%), ³³S (0.75%±0.02%), ³⁴S (4.25%±0.24%), and ³⁶S (0.01%±0.01%). Other than ³⁵S, with a half-life of 87 days and formed in cosmic ray spallation of ⁴⁰Ar, the radioactive isotopes of sulfur have half-lives less than 3 hours. 

When sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the δS-34 values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The δC-13 and δS-34 of coexisting carbonate minerals and sulfides can be used to determine the pH and oxygen fugacity of the ore-bearing fluid during ore formation. 

In most forest ecosystems, sulfate is derived mostly from the atmosphere; weathering of ore minerals and evaporites contribute some sulfur. Sulfur with a distinctive isotopic composition has been used to identify pollution sources, and enriched sulfur has been added as a tracer in hydrologic studies. Differences in the natural abundances can be used in systems where there is sufficient variation in the ³⁴S of ecosystem components. Rocky Mountain lakes thought to be dominated by atmospheric sources of sulfate have been found to have different δ³⁴S values from lakes believed to be dominated by watershed sources of sulfate.

Natural occurrence

Sulfur vat from which railroad cars are loaded, Freeport Sulphur Co., Hoskins Mound, Texas (1943)

 

Most of the yellow and orange hues of Io are due to elemental sulfur and sulfur compounds deposited by active volcanoes.

 

A man carrying sulfur blocks from Kawah Ijen, a volcano in East Java, Indonesia, 2009

 

³²S is created inside massive stars, at a depth where the temperature exceeds 2.5×10⁹ K, by the fusion of one nucleus of silicon plus one nucleus of helium. As this nuclear reaction is part of the alpha process that produces elements in abundance, sulfur is the 10th most common element in the universe. 

Sulfur, usually as sulfide, is present in many types of meteorites. Ordinary chondrites contain on average 2.1% sulfur, and carbonaceous chondrites may contain as much as 6.6%. It is normally present as troilite (FeS), but there are exceptions, with carbonaceous chondrites containing free sulfur, sulfates and other sulfur compounds. The distinctive colors of Jupiter's volcanic moon Io are attributed to various forms of molten, solid, and gaseous sulfur.

Sulfur occurs in fumaroles such as this one in Vulcano, Italy

 

It is the fifth most common element by mass in the Earth. Elemental sulfur can be found near hot springs and volcanic regions in many parts of the world, especially along the Pacific Ring of Fire; such volcanic deposits are currently mined in Indonesia, Chile, and Japan. These deposits are polycrystalline, with the largest documented single crystal measuring 22×16×11 cm. Historically, Sicily was a major source of sulfur in the Industrial Revolution.

Native sulfur is synthesized by anaerobic bacteria acting on sulfate minerals such as gypsum in salt domes. Significant deposits in salt domes occur along the coast of the Gulf of Mexico, and in evaporites in eastern Europe and western Asia. Native sulfur may be produced by geological processes alone. Fossil-based sulfur deposits from salt domes were until recently the basis for commercial production in the United States, Russia, Turkmenistan, and Ukraine. Currently, commercial production is still carried out in the Osiek mine in Poland. Such sources are now of secondary commercial importance, and most are no longer worked.

Common naturally occurring sulfur compounds include the sulfide minerals, such as pyrite (iron sulfide), cinnabar (mercury sulfide), galena (lead sulfide), sphalerite (zinc sulfide) and stibnite (antimony sulfide); and the sulfates, such as gypsum (calcium sulfate), alunite (potassium aluminium sulfate), and barite (barium sulfate). On Earth, just as upon Jupiter's moon Io, elemental sulfur occurs naturally in volcanic emissions, including emissions from hydrothermal vents.

Compounds

Common oxidation states of sulfur range from −2 to +6. Sulfur forms stable compounds with all elements except the noble gases.

Sulfur polycations

Sulfur polycations, S₈²⁺, S₄²⁺ and S₁₆²⁺ are produced when sulfur is reacted with mild oxidizing agents in a strongly acidic solution. The colored solutions produced by dissolving sulfur in oleum were first reported as early as 1804 by C.F. Bucholz, but the cause of the color and the structure of the polycations involved was only determined in the late 1960s. S₈²⁺ is deep blue, S₄²⁺ is yellow and S₁₆²⁺ is red.

Sulfides

Treatment of sulfur with hydrogen gives hydrogen sulfide. When dissolved in water, hydrogen sulfide is mildly acidic:

H₂S ⇌ HS⁻ + H⁺

Hydrogen sulfide gas and the hydrosulfide anion are extremely toxic to mammals, due to their inhibition of the oxygen-carrying capacity of hemoglobin and certain cytochromes in a manner analogous to cyanide and azide. 

Reduction of elemental sulfur gives polysulfides, which consist of chains of sulfur atoms terminated with S⁻ centers:

2 Na + S₈ → Na₂S₈

This reaction highlights a distinctive property of sulfur: its ability to catenate (bind to itself by formation of chains). Protonation of these polysulfide anions produces the polysulfanes, H₂S_x where x = 2, 3, and 4. Ultimately, reduction of sulfur produces sulfide salts:

16 Na + S₈ → 8 Na₂S

The interconversion of these species is exploited in the sodium-sulfur battery. 

The radical anion S₃⁻ gives the blue color of the mineral lapis lazuli. 

Lapis lazuli owes its blue color to a trisulfur radical anion (S⁻
₃)

 

Two parallel sulfur chains grown inside a single-wall carbon nanotube (CNT, a). Zig-zag (b) and straight (c) S chains inside double-wall CNTs

Oxides, oxoacids and oxoanions

The principal sulfur oxides are obtained by burning sulfur:

S + O₂ → SO₂ (sulfur dioxide)

2 SO₂ + O₂ → 2 SO₃ (sulfur trioxide)

Multiple sulfur oxides are known; the sulfur-rich oxides include sulfur monoxide, disulfur monoxide, disulfur dioxides, and higher oxides containing peroxo groups. 

Sulfur forms sulfur oxoacids, some of which cannot be isolated and are only known through the salts. Sulfur dioxide and sulfites (SO²⁻
₃) are related to the unstable sulfurous acid (H₂SO₃). Sulfur trioxide and sulfates (SO²⁻
₄) are related to sulfuric acid (H₂SO₄). Sulfuric acid and SO₃ combine to give oleum, a solution of pyrosulfuric acid (H₂S₂O₇) in sulfuric acid.

Thiosulfate salts (S
₂O²⁻
₃), sometimes referred as "hyposulfites", used in photographic fixing (hypo) and as reducing agents, feature sulfur in two oxidation states. Sodium dithionite (Na
₂S
₂O
₄), contains the more highly reducing dithionite anion (S
₂O²⁻
₄).

Halides and oxyhalides

Several sulfur halides are important to modern industry. Sulfur hexafluoride is a dense gas used as an insulator gas in high voltage transformers; it is also a nonreactive and nontoxic propellant for pressurized containers. Sulfur tetrafluoride is a rarely used organic reagent that is highly toxic. Sulfur dichloride and disulfur dichloride are important industrial chemicals. Sulfuryl chloride and chlorosulfuric acid are derivatives of sulfuric acid; thionyl chloride (SOCl₂) is a common reagent in organic synthesis.

Pnictides

An important S–N compound is the cage tetrasulfur tetranitride (S₄N₄). Heating this compound gives polymeric sulfur nitride ((SN)_x), which has metallic properties even though it does not contain any metal atoms. Thiocyanates contain the SCN⁻ group. Oxidation of thiocyanate gives thiocyanogen, (SCN)₂ with the connectivity NCS-SCN. Phosphorus sulfides are numerous, the most important commercially being the cages P₄S₁₀ and P₄S₃.

Metal sulfides

The principal ores of copper, zinc, nickel, cobalt, molybdenum, and other metals are sulfides. These materials tend to be dark-colored semiconductors that are not readily attacked by water or even many acids. They are formed, both geochemically and in the laboratory, by the reaction of hydrogen sulfide with metal salts. The mineral galena (PbS) was the first demonstrated semiconductor and was used as a signal rectifier in the cat's whiskers of early crystal radios. The iron sulfide called pyrite, the so-called "fool's gold", has the formula FeS₂. Processing these ores, usually by roasting, is costly and environmentally hazardous. Sulfur corrodes many metals through tarnishing.

Organic compounds

• Illustrative organosulfur compounds
• 

  Allicin, the active ingredient in garlic
  
• 

  (R)-cysteine, an amino acid containing a thiol group
  
• 

  Methionine, an amino acid containing a thioether
  
• 

  Diphenyl disulfide, a representative disulfide
  
• 

  Perfluorooctanesulfonic acid, a controversial surfactant
  
• 

  Dibenzothiophene, a component of crude oil
  
• 

  Penicillin, an antibiotic where "R" is the variable group
  

Some of the main classes of sulfur-containing organic compounds include the following:

• Thiols or mercaptans (so called because they capture mercury as chelators) are the sulfur analogs of alcohols; treatment of thiols with base gives thiolate ions.
• Thioethers are the sulfur analogs of ethers.
• Sulfonium ions have three groups attached to a cationic sulfur center. Dimethylsulfoniopropionate (DMSP) is one such compound, important in the marine organic sulfur cycle.
• Sulfoxides and sulfones are thioethers with one and two oxygen atoms attached to the sulfur atom, respectively. The simplest sulfoxide, dimethyl sulfoxide, is a common solvent; a common sulfone is sulfolane.
• Sulfonic acids are used in many detergents.

Compounds with carbon-sulfur multiple bonds are uncommon, an exception being carbon disulfide, a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymer rayon and many organosulfur compounds. Unlike carbon monoxide, carbon monosulfide is stable only as an extremely dilute gas, found between solar systems.

Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as the odorant in domestic natural gas, garlic odor, and skunk spray. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing monoterpenoid (grapefruit mercaptan) in small concentrations is the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations. Sulfur mustard, a potent vesicant, was used in World War I as a disabling agent.

Sulfur-sulfur bonds are a structural component used to stiffen rubber, similar to the disulfide bridges that rigidify proteins (see biological below). In the most common type of industrial "curing" or hardening and strengthening of natural rubber, elemental sulfur is heated with the rubber to the point that chemical reactions form disulfide bridges between isoprene units of the polymer. This process, patented in 1843, made rubber a major industrial product, especially in automobile tires. Because of the heat and sulfur, the process was named vulcanization, after the Roman god of the forge and volcanism.

History

Antiquity

Pharmaceutical container for sulfur from the first half of the 20th century. From the Museo del Objeto del Objeto collection

 

Being abundantly available in native form, sulfur was known in ancient times and is referred to in the Torah (Genesis). English translations of the Bible commonly referred to burning sulfur as "brimstone", giving rise to the term "fire-and-brimstone" sermons, in which listeners are reminded of the fate of eternal damnation that await the unbelieving and unrepentant. It is from this part of the Bible that Hell is implied to "smell of sulfur" (likely due to its association with volcanic activity). According to the Ebers Papyrus, a sulfur ointment was used in ancient Egypt to treat granular eyelids. Sulfur was used for fumigation in preclassical Greece; this is mentioned in the Odyssey. Pliny the Elder discusses sulfur in book 35 of his Natural History, saying that its best-known source is the island of Melos. He mentions its use for fumigation, medicine, and bleaching cloth.

A natural form of sulfur known as shiliuhuang (石硫黄) was known in China since the 6th century BC and found in Hanzhong. By the 3rd century, the Chinese discovered that sulfur could be extracted from pyrite. Chinese Daoists were interested in sulfur's flammability and its reactivity with certain metals, yet its earliest practical uses were found in traditional Chinese medicine. A Song dynasty military treatise of 1044 AD described different formulas for Chinese black powder, which is a mixture of potassium nitrate (KNO
₃), charcoal, and sulfur. It remains an ingredient of black gunpowder. 

Indian alchemists, practitioners of "the science of mercury" (sanskrit rasaśāstra, रसशास्त्र), wrote extensively about the use of sulfur in alchemical operations with mercury, from the eighth century AD onwards. In the rasaśāstra tradition, sulfur is called "the smelly" (sanskrit gandhaka, गन्धक).

Various alchemical symbols for sulfur

 

Early European alchemists gave sulfur a unique alchemical symbol, a triangle at the top of a cross. In traditional skin treatment, elemental sulfur was used (mainly in creams) to alleviate such conditions as scabies, ringworm, psoriasis, eczema, and acne. The mechanism of action is unknown—though elemental sulfur does oxidize slowly to sulfurous acid, which is (through the action of sulfite) a mild reducing and antibacterial agent.

Modern times

Sicilian kiln used to obtain sulfur from volcanic rock

In 1777, Antoine Lavoisier helped convince the scientific community that sulfur was an element, not a compound.

Sulfur deposits in Sicily were the dominant source for more than a century. By the late 18th century, about 2,000 tonnes per year of sulfur were imported into Marseilles, France, for the production of sulfuric acid for use in the Leblanc process. In industrializing Britain, with the repeal of tariffs on salt in 1824, demand for sulfur from Sicily surged upward. The increasing British control and exploitation of the mining, refining, and transportation of the sulfur, coupled with the failure of this lucrative export to transform Sicily's backward and impoverished economy, led to the 'Sulfur Crisis' of 1840, when King Ferdinand II gave a monopoly of the sulfur industry to a French firm, violating an earlier 1816 trade agreement with Britain. A peaceful solution was eventually negotiated by France.

In 1867, elemental sulfur was discovered in underground deposits in Louisiana and Texas. The highly successful Frasch process was developed to extract this resource.

In the late 18th century, furniture makers used molten sulfur to produce decorative inlays in their craft. Because of the sulfur dioxide produced during the process of melting sulfur, the craft of sulfur inlays was soon abandoned. Molten sulfur is sometimes still used for setting steel bolts into drilled concrete holes where high shock resistance is desired for floor-mounted equipment attachment points. Pure powdered sulfur was used as a medicinal tonic and laxative. With the advent of the contact process, the majority of sulfur today is used to make sulfuric acid for a wide range of uses, particularly fertilizer.

Spelling and etymology

Sulfur is derived from the Latin word sulpur, which was Hellenized to sulphur. The spelling sulfur appears toward the end of the Classical period. (The true Greek word for sulfur, θεῖον, is the source of the international chemical prefix thio-.) In 12th-century Anglo-French, it was sulfre; in the 14th century the Latin -ph- was restored, for sulphre; and by the 15th century the full Latin spelling was restored, for sulfur, sulphur. The parallel f~ph spellings continued in Britain until the 19th century, when the word was standardized as sulphur. Sulfur was the form chosen in the United States, whereas Canada uses both. The IUPAC adopted the spelling sulfur in 1990, as did the Nomenclature Committee of the Royal Society of Chemistry in 1992, restoring the spelling sulfur to Britain. Oxford Dictionaries note that "in chemistry and other technical uses ... the -f- spelling is now the standard form for this and related words in British as well as US contexts, and is increasingly used in general contexts as well."

Production

Traditional sulfur mining at Ijen Volcano, East Java, Indonesia. This image shows the dangerous and rugged conditions the miners face, including toxic smoke and high drops, as well as their lack of protective equipment. The pipes over which they are standing are for condensing sulfur vapors.

 

Sulfur may be found by itself and historically was usually obtained in this form; pyrite has also been a source of sulfur. In volcanic regions in Sicily, in ancient times, it was found on the surface of the Earth, and the "Sicilian process" was used: sulfur deposits were piled and stacked in brick kilns built on sloping hillsides, with airspaces between them. Then, some sulfur was pulverized, spread over the stacked ore and ignited, causing the free sulfur to melt down the hills. Eventually the surface-borne deposits played out, and miners excavated veins that ultimately dotted the Sicilian landscape with labyrinthine mines. Mining was unmechanized and labor-intensive, with pickmen freeing the ore from the rock, and mine-boys or carusi carrying baskets of ore to the surface, often through a mile or more of tunnels. Once the ore was at the surface, it was reduced and extracted in smelting ovens. The conditions in Sicilian sulfur mines were horrific, prompting Booker T. Washington to write "I am not prepared just now to say to what extent I believe in a physical hell in the next world, but a sulphur mine in Sicily is about the nearest thing to hell that I expect to see in this life."

Elemental sulfur was extracted from salt domes (in which it sometimes occurs in nearly pure form) until the late 20th century. Sulfur is now produced as a side product of other industrial processes such as in oil refining, in which sulfur is undesired. As a mineral, native sulfur under salt domes is thought to be a fossil mineral resource, produced by the action of ancient bacteria on sulfate deposits. It was removed from such salt-dome mines mainly by the Frasch process. In this method, superheated water was pumped into a native sulfur deposit to melt the sulfur, and then compressed air returned the 99.5% pure melted product to the surface. Throughout the 20th century this procedure produced elemental sulfur that required no further purification. Due to a limited number of such sulfur deposits and the high cost of working them, this process for mining sulfur has not been employed in a major way anywhere in the world since 2002.

Sulfur recovered from hydrocarbons in Alberta, stockpiled for shipment in North Vancouver, British Columbia

 

Today, sulfur is produced from petroleum, natural gas, and related fossil resources, from which it is obtained mainly as hydrogen sulfide. Organosulfur compounds, undesirable impurities in petroleum, may be upgraded by subjecting them to hydrodesulfurization, which cleaves the C–S bonds:

R-S-R + 2 H₂ → 2 RH + H₂S

The resulting hydrogen sulfide from this process, and also as it occurs in natural gas, is converted into elemental sulfur by the Claus process. This process entails oxidation of some hydrogen sulfide to sulfur dioxide and then the comproportionation of the two:

3 O₂ + 2 H₂S → 2 SO₂ + 2 H₂O

SO₂ + 2 H₂S → 3 S + 2 H₂O

Production and price (US market) of elemental sulfur

 

Owing to the high sulfur content of the Athabasca Oil Sands, stockpiles of elemental sulfur from this process now exist throughout Alberta, Canada. Another way of storing sulfur is as a binder for concrete, the resulting product having many desirable properties. Sulfur is still mined from surface deposits in poorer nations with volcanoes, such as Indonesia, and worker conditions have not improved much since Booker T. Washington's days.

The world production of sulfur in 2011 amounted to 69 million tonnes (Mt), with more than 15 countries contributing more than 1 Mt each. Countries producing more than 5 Mt are China (9.6), US (8.8), Canada (7.1) and Russia (7.1). Production has been slowly increasing from 1900 to 2010; the price was unstable in the 1980s and around 2010.

Applications

Sulfuric acid

Elemental sulfur is used mainly as a precursor to other chemicals. Approximately 85% (1989) is converted to sulfuric acid (H₂SO₄):

2 S + 3 O₂ + 2 H₂O → 2 H₂SO₄

 

Sulfuric acid production in 2000

In 2010, the United States produced more sulfuric acid than any other inorganic industrial chemical. The principal use for the acid is the extraction of phosphate ores for the production of fertilizer manufacturing. Other applications of sulfuric acid include oil refining, wastewater processing, and mineral extraction.

Other important sulfur chemistry

Sulfur reacts directly with methane to give carbon disulfide, used to manufacture cellophane and rayon. One of the uses of elemental sulfur is in vulcanization of rubber, where polysulfide chains crosslink organic polymers. Large quantities of sulfites are used to bleach paper and to preserve dried fruit. Many surfactants and detergents (e.g. sodium lauryl sulfate) are sulfate derivatives. Calcium sulfate, gypsum, (CaSO₄·2H₂O) is mined on the scale of 100 million tonnes each year for use in Portland cement and fertilizers.

When silver-based photography was widespread, sodium and ammonium thiosulfate were widely used as "fixing agents". Sulfur is a component of gunpowder ("black powder").

Fertilizer

Sulfur is increasingly used as a component of fertilizers. The most important form of sulfur for fertilizer is the mineral calcium sulfate. Elemental sulfur is hydrophobic (not soluble in water) and cannot be used directly by plants. Over time, soil bacteria can convert it to soluble derivatives, which can then be used by plants. Sulfur improves the efficiency of other essential plant nutrients, particularly nitrogen and phosphorus. Biologically produced sulfur particles are naturally hydrophilic due to a biopolymer coating and are easier to disperse over the land in a spray of diluted slurry, resulting in a faster uptake. 

The botanical requirement for sulfur equals or exceeds the requirement for phosphorus. It is an essential nutrient for plant growth, root nodule formation of legumes, and immunity and defense systems. Sulfur deficiency has become widespread in many countries in Europe. Because atmospheric inputs of sulfur continue to decrease, the deficit in the sulfur input/output is likely to increase unless sulfur fertilizers are used.

Fine chemicals

A molecular model of the pesticide malathion

 

Organosulfur compounds are used in pharmaceuticals, dyestuffs, and agrochemicals. Many drugs contain sulfur, early examples being antibacterial sulfonamides, known as sulfa drugs. Sulfur is a part of many bacterial defense molecules. Most β-lactam antibiotics, including the penicillins, cephalosporins and monolactams contain sulfur.

Magnesium sulfate, known as Epsom salts when in hydrated crystal form, can be used as a laxative, a bath additive, an exfoliant, magnesium supplement for plants, or (when in dehydrated form) as a desiccant.

Fungicide and pesticide

Sulfur candle originally sold for home fumigation

 

Elemental sulfur is one of the oldest fungicides and pesticides. "Dusting sulfur", elemental sulfur in powdered form, is a common fungicide for grapes, strawberry, many vegetables and several other crops. It has a good efficacy against a wide range of powdery mildew diseases as well as black spot. In organic production, sulfur is the most important fungicide. It is the only fungicide used in organically farmed apple production against the main disease apple scab under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophilic characteristics) can also be used for these applications. 

Standard-formulation dusting sulfur is applied to crops with a sulfur duster or from a dusting plane. Wettable sulfur is the commercial name for dusting sulfur formulated with additional ingredients to make it water miscible. It has similar applications and is used as a fungicide against mildew and other mold-related problems with plants and soil.

Elemental sulfur powder is used as an "organic" (i.e., "green") insecticide (actually an acaricide) against ticks and mites. A common method of application is dusting the clothing or limbs with sulfur powder. 

A diluted solution of lime sulfur (made by combining calcium hydroxide with elemental sulfur in water) is used as a dip for pets to destroy ringworm (fungus), mange, and other dermatoses and parasites. 

Sulfur candles of almost pure sulfur were burned to fumigate structures and wine barrels, but are now considered too toxic for residences.

Bactericide in winemaking and food preservation

Small amounts of sulfur dioxide gas addition (or equivalent potassium metabisulfite addition) to fermented wine to produce traces of sulfurous acid (produced when SO₂ reacts with water) and its sulfite salts in the mixture, has been called "the most powerful tool in winemaking". After the yeast-fermentation stage in winemaking, sulfites absorb oxygen and inhibit aerobic bacterial growth that otherwise would turn ethanol into acetic acid, souring the wine. Without this preservative step, indefinite refrigeration of the product before consumption is usually required. Similar methods go back into antiquity but modern historical mentions of the practice go to the fifteenth century. The practice is used by large industrial wine producers and small organic wine producers alike. 

Sulfur dioxide and various sulfites have been used for their antioxidant antibacterial preservative properties in many other parts of the food industry. The practice has declined since reports of an allergy-like reaction of some persons to sulfites in foods.

Pharmaceuticals

Octasulfur

Clinical data
AHFS/Drugs.com	Multum Consumer Information
Routes of administration	Topical, rarely oral
ATC code	D10AB02 (WHO)
Legal status
Legal status	US: OTC In general: Over-the-counter (OTC)
Identifiers
IUPAC name[show]
CAS Number	10544-50-0
PubChem CID	66348
ChemSpider	59726
ChEBI	CHEBI:29385
ECHA InfoCard	100.028.839
Chemical and physical data
Formula	S8
Molar mass	256.52 g/mol g·mol−1

Sulfur (specifically octasulfur, S₈) is used in pharmaceutical skin preparations for the treatment of acne and other conditions. It acts as a keratolytic agent and also kills bacteria, fungi, scabies mites and other parasites. Precipitated sulfur and colloidal sulfur are used, in form of lotions, creams, powders, soaps, and bath additives, for the treatment of acne vulgaris, acne rosacea, and seborrhoeic dermatitis.

Common adverse effects include irritation of the skin at the application site, such as dryness, stinging, itching and peeling.

Mechanism of action

Sulfur is converted to hydrogen sulfide (H₂S) through reduction, partly by bacteria. H₂S kills bacteria (possibly including Propionibacterium acnes which plays a role in acne,) fungi, and parasites such as scabies mites. Sulfur's keratolytic action is also mediated by H₂S; in this case, the hydrogen sulfide is produced by direct interaction with the target keratinocytes themselves.

Furniture

Sulfur can be used to create decorative inlays in wooden furniture. After a design has been cut into the wood, molten sulfur is poured in and then scraped away so it is flush. Sulfur inlays were particularly popular in the late 18th and early 19th centuries, notably among Pennsylvania German cabinetmakers. The practice soon died out, as less toxic and flammable substances were substituted. However, some modern craftsmen have occasionally revived the technique in the creation of replica pieces.

Biological role

Protein and organic cofactors

Sulfur is an essential component of all living cells. It is either the seventh or eighth most abundant element in the human body by weight, about equal in abundance to potassium, and slightly greater than sodium and chlorine. A 70 kg (150 lb) human body contains about 140 grams of sulfur. 

In plants and animals, the amino acids cysteine and methionine contain most of the sulfur, and the element is present in all polypeptides, proteins, and enzymes that contain these amino acids. In humans, methionine is an essential amino acid that must be ingested. However, save for the vitamins biotin and thiamine, cysteine and all sulfur-containing compounds in the human body can be synthesized from methionine. The enzyme sulfite oxidase is needed for the metabolism of methionine and cysteine in humans and animals. 

Disulfide bonds (S-S bonds) between cysteine residues in peptide chains are very important in protein assembly and structure. These covalent bonds between peptide chains confer extra toughness and rigidity. For example, the high strength of feathers and hair is due in part to the high content of S-S bonds with cysteine and sulfur. Eggs are high in sulfur to nourish feather formation in chicks, and the characteristic odor of rotting eggs is due to hydrogen sulfide. The high disulfide bond content of hair and feathers contributes to their indigestibility and to their characteristic disagreeable odor when burned. 

Homocysteine and taurine are other sulfur-containing acids that are similar in structure, but not coded by DNA, and are not part of the primary structure of proteins. Many important cellular enzymes use prosthetic groups ending with -SH moieties to handle reactions involving acyl-containing biochemicals: two common examples from basic metabolism are coenzyme A and alpha-lipoic acid. Two of the 13 classical vitamins, biotin and thiamine, contain sulfur, with the latter being named for its sulfur content. 

In intracellular chemistry, sulfur operates as a carrier of reducing hydrogen and its electrons for cellular repair of oxidation. Reduced glutathione, a sulfur-containing tripeptide, is a reducing agent through its sulfhydryl (-SH) moiety derived from cysteine. The thioredoxins, a class of small proteins essential to all known life, use neighboring pairs of reduced cysteines to work as general protein reducing agents, with similar effect. 

Methanogenesis, the route to most of the world's methane, is a multistep biochemical transformation of carbon dioxide. This conversion requires several organosulfur cofactors. These include coenzyme M, CH₃SCH₂CH₂SO₃⁻, the immediate precursor to methane.

Metalloproteins and inorganic cofactors

Inorganic sulfur forms a part of iron–sulfur clusters as well as many copper, nickel, and iron proteins. Most pervasive are the ferrodoxins, which serve as electron shuttles in cells. In bacteria, the important nitrogenase enzymes contains an Fe–Mo–S cluster and is a catalyst that performs the important function of nitrogen fixation, converting atmospheric nitrogen to ammonia that can be used by microorganisms and plants to make proteins, DNA, RNA, alkaloids, and the other organic nitrogen compounds necessary for life.

Sulfur metabolism and the sulfur cycle

The sulfur cycle was the first of the biogeochemical cycles to be discovered. In the 1880s, while studying Beggiatoa (a bacterium living in a sulfur rich environment), Sergei Winogradsky found that it oxidized hydrogen sulfide (H₂S) as an energy source, forming intracellular sulfur droplets. Winogradsky referred to this form of metabolism as inorgoxidation (oxidation of inorganic compounds). He continued to study it together with Selman Waksman until the 1950s.

Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur, sulfite, thiosulfate, and various polythionates (e.g., tetrathionate). They depend on enzymes such as sulfur oxygenase and sulfite oxidase to oxidize sulfur to sulfate. Some lithotrophs can even use the energy contained in sulfur compounds to produce sugars, a process known as chemosynthesis. Some bacteria and archaea use hydrogen sulfide in place of water as the electron donor in chemosynthesis, a process similar to photosynthesis that produces sugars and utilizes oxygen as the electron acceptor. The photosynthetic green sulfur bacteria and purple sulfur bacteria and some lithotrophs use elemental oxygen to carry out such oxidization of hydrogen sulfide to produce elemental sulfur (S⁰), oxidation state = 0. Primitive bacteria that live around deep ocean volcanic vents oxidize hydrogen sulfide in this way with oxygen; the giant tube worm is an example of a large organism that uses hydrogen sulfide (via bacteria) as food to be oxidized. 

The so-called sulfate-reducing bacteria, by contrast, "breathe sulfate" instead of oxygen. They use organic compounds or molecular hydrogen as the energy source. They use sulfur as the electron acceptor, and reduce various oxidized sulfur compounds back into sulfide, often into hydrogen sulfide. They can grow on other partially oxidized sulfur compounds (e.g. thiosulfates, thionates, polysulfides, sulfites). The hydrogen sulfide produced by these bacteria is responsible for some of the smell of intestinal gases (flatus) and decomposition products.

Sulfur is absorbed by plants roots from soil as sulfate and transported as a phosphate ester. Sulfate is reduced to sulfide via sulfite before it is incorporated into cysteine and other organosulfur compounds.

SO₄²⁻ → SO₃²⁻ → H₂S → cysteine → methionine

Precautions

Sulfur

Hazards
GHS pictograms
GHS signal word	Warning
GHS hazard statements	H315
NFPA 704	1 2 2

Effect of acid rain on a forest, Jizera Mountains, Czech Republic

 

Elemental sulfur is non-toxic, as are most of the soluble sulfate salts, such as Epsom salts. Soluble sulfate salts are poorly absorbed and laxative. When injected parenterally, they are freely filtered by the kidneys and eliminated with very little toxicity in multi-gram amounts. 

When sulfur burns in air, it produces sulfur dioxide. In water, this gas produces sulfurous acid and sulfites; sulfites are antioxidants that inhibit growth of aerobic bacteria and a useful food additive in small amounts. At high concentrations these acids harm the lungs, eyes or other tissues. In organisms without lungs such as insects or plants, sulfite in high concentration prevents respiration.

Sulfur trioxide (made by catalysis from sulfur dioxide) and sulfuric acid are similarly highly acidic and corrosive in the presence of water. Sulfuric acid is a strong dehydrating agent that can strip available water molecules and water components from sugar and organic tissue.

The burning of coal and/or petroleum by industry and power plants generates sulfur dioxide (SO₂) that reacts with atmospheric water and oxygen to produce sulfuric acid (H₂SO₄) and sulfurous acid (H₂SO₃). These acids are components of acid rain, lowering the pH of soil and freshwater bodies, sometimes resulting in substantial damage to the environment and chemical weathering of statues and structures. Fuel standards increasingly require that fuel producers extract sulfur from fossil fuels to prevent acid rain formation. This extracted and refined sulfur represents a large portion of sulfur production. In coal-fired power plants, flue gases are sometimes purified. More modern power plants that use synthesis gas extract the sulfur before they burn the gas.

Hydrogen sulfide is as toxic as hydrogen cyanide, and kills by the same mechanism (inhibition of the respiratory enzyme cytochrome oxidase), though hydrogen sulfide is less likely to cause surprise poisonings from small inhaled amounts because of its disagreeable odor. Hydrogen sulfide quickly deadens the sense of smell and a victim may breathe increasing quantities without noticing the increase until severe symptoms cause death. Dissolved sulfide and hydrosulfide salts are toxic by the same mechanism.