Sulfuric acid (updated)

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Sulfuric acid

Space-filling model Ball-and-stick model
Names
IUPAC name Sulfuric acid
Other names Oil of vitriol
Identifiers
CAS Number	7664-93-9
3D model (JSmol)	Interactive image
ChEBI	CHEBI:26836
ChEMBL	ChEMBL572964
ChemSpider	1086
ECHA InfoCard	100.028.763
EC Number	231-639-5
E number	E513
Gmelin Reference	2122
KEGG	D05963
PubChem CID	1118
RTECS number	WS5600000
UNII	O40UQP6WCF
UN number	1830
Properties
Chemical formula	H 2SO 4
Molar mass	98.079 g/mol
Appearance	Clear, colorless liquid
Odor	odorless
Density	1.8302 g/cm3, liquid
Melting point	10.31 °C (50.56 °F; 283.46 K)
Boiling point	337 °C (639 °F; 610 K) When sulfuric acid is above 300 °C (572 °F), it will decompose slowly
Solubility in water	miscible, exothermic
Vapor pressure	0.001 mmHg (20 °C)
Acidity (pKa)	−3, 1.99
Conjugate base	Hydrogen sulfate
Viscosity	26.7 cP (20 °C)
Thermochemistry
Std molar entropy (So298)	157 J·mol−1·K−1
Std enthalpy of formation (ΔfHo298)	−814 kJ·mol−1
Hazards
Safety data sheet	External MSDS
GHS pictograms
GHS signal word	Danger
GHS hazard statements	H314
GHS precautionary statements	P260, P264, P280, P301+330+331, P303+361+353, P363, P304+340, P305+351+338, P310, P321, P405, P501
NFPA 704	0 3 2 W OX
Flash point	Non-flammable
Threshold limit value	15 mg/m3 (IDLH), 1 mg/m3 (TWA), 2 mg/m3 (STEL)
Lethal dose or concentration (LD, LC):
LD50 (median dose)	2140 mg/kg (rat, oral)
LC50 (median concentration)	50 mg/m3 (guinea pig, 8 hr) 510 mg/m3 (rat, 2 hr) 320 mg/m3 (mouse, 2 hr) 18 mg/m3 (guinea pig)
LCLo (lowest published)	87 mg/m3 (guinea pig, 2.75 hr)
US health exposure limits (NIOSH):
PEL (Permissible)	TWA 1 mg/m3
REL (Recommended)	TWA 1 mg/m3
IDLH (Immediate danger)	15 mg/m3
Related compounds
Related strong acids	Selenic acid Hydrochloric acid Nitric acid Chromic acid
Related compounds	Sulfurous acid Peroxymonosulfuric acid Sulfur trioxide Oleum
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Sulfuric acid (alternative spelling sulphuric acid), also known as vitriol, is a mineral acid composed of the elements sulfur, oxygen and hydrogen, with molecular formula H₂SO₄. It is a colorless, odorless, and syrupy liquid that is soluble in water, in a reaction that is highly exothermic.

Its corrosiveness can be mainly ascribed to its strong acidic nature, and, if at a high concentration, its dehydrating and oxidizing properties. It is also hygroscopic, readily absorbing water vapor from the air. Upon contact, sulfuric acid can cause severe chemical burns and even secondary thermal burns; it is very dangerous even at moderate concentrations.

Sulfuric acid is a very important commodity chemical, and indeed, a nation's sulfuric acid production is a good indicator of its industrial strength. It is widely produced with different methods, such as contact process, wet sulfuric acid process, lead chamber process and some other methods.

The most common use of sulfuric acid is for fertilizer manufacture. It is also a central substance in the chemical industry. Principal uses include fertilizer manufacturing (and other mineral processing), oil refining, wastewater processing, and chemical synthesis. It has a wide range of end applications including in domestic acidic drain cleaners, as an electrolyte in lead-acid batteries and in various cleaning agents.

Physical properties

Grades of sulfuric acid

Although nearly 100% sulfuric acid can be made, the subsequent loss of SO
₃ at the boiling point brings the concentration to 98.3% acid. The 98.3% grade is more stable in storage, and is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes. Some common concentrations are:

Mass fraction H2SO4	Density (kg/L)	Concentration (mol/L)	Common name
l.t. 29%	1.00-1.25	l.t. 4.2	diluted sulfuric acid
29–32%	1.25–1.28	4.2–5.0	battery acid (used in lead–acid batteries)
62–70%	1.52–1.60	9.6–11.5	chamber acid fertilizer acid
78–80%	1.70–1.73	13.5–14.0	tower acid Glover acid
98.3%	1.84	18.4	concentrated sulfuric acid

"Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process, chamber acid being the acid produced in the lead chamber itself (less than 70% to avoid contamination with nitrosylsulfuric acid) and tower acid being the acid recovered from the bottom of the Glover tower. They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10M" sulfuric acid (the modern equivalent of chamber acid, used in many titrations) is prepared by slowly adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C (176 °F) or higher.

Sulfuric acid reacts with its anhydride, SO
₃, to form H
₂S
₂O
₇, called pyrosulfuric acid, fuming sulfuric acid, Disulfuric acid or oleum or, less commonly, Nordhausen acid. Concentrations of oleum are either expressed in terms of % SO
₃ (called % oleum) or as % H
₂SO
₄ (the amount made if H
₂O were added); common concentrations are 40% oleum (109% H
₂SO
₄) and 65% oleum (114.6% H
₂SO
₄). Pure H
₂S
₂O
₇ is a solid with melting point of 36 °C. 

Pure sulfuric acid has a vapor pressure of less than 0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, and 98% sulfuric acid has a less than 1 mmHg vapor pressure at 40 °C.

Pure sulfuric acid is a viscous clear liquid, like oil, and this explains the old name of the acid ('oil of vitriol'). 

Commercial sulfuric acid is sold in several different purity grades. Technical grade H
₂SO
₄ is impure and often colored, but is suitable for making fertilizer. Pure grades, such as United States Pharmacopeia (USP) grade, are used for making pharmaceuticals and dyestuffs. Analytical grades are also available. 

Nine hydrates are known, but four of them were confirmed to be tetrahydrate (H₂SO₄·4H₂O), hemihexahydrate (H₂SO₄·​6 ¹⁄₂H₂O) and octahydrate (H₂SO₄·8H₂O).

Polarity and conductivity

Equilibrium of anhydrous sulfuric acid

Species	mMol/kg
HSO− 4	15.0
H 3SO+ 4	11.3
H 3O+	8.0
HS 2O− 7	4.4
H 2S 2O 7	3.6
H 2O	0.1

Anhydrous H
₂SO
₄ is a very polar liquid, having a dielectric constant of around 100. It has a high electrical conductivity, caused by dissociation through protonating itself, a process known as autoprotolysis.

2 H
₂SO
₄ ⇌ H
₃SO⁺
₄ + HSO⁻
₄

The equilibrium constant for the autoprotolysis is

K_ap (25 °C) = [H
₃SO⁺
₄][HSO⁻
₄] = 2.7×10⁻⁴

The comparable equilibrium constant for water, K_w is 10⁻¹⁴, a factor of 10¹⁰ (10 billion) smaller.

In spite of the viscosity of the acid, the effective conductivities of the H
₃SO⁺
₄ and HSO⁻
₄ ions are high due to an intramolecular proton-switch mechanism (analogous to the Grotthuss mechanism in water), making sulfuric acid a good conductor of electricity. It is also an excellent solvent for many reactions.

Chemical properties

Reaction with water and dehydrating property

Drops of concentrated sulfuric acid rapidly decompose a piece of cotton towel by dehydration.

Because the hydration reaction of sulfuric acid is highly exothermic, dilution should always be performed by adding the acid to the water rather than the water to the acid. Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the acid is the limiting reagent. This reaction is best thought of as the formation of hydronium ions:

H
₂SO
₄ + H
₂O → H
₃O⁺ + HSO⁻
₄   K_a1 = 2.4×10⁶ (strong acid)

HSO⁻
₄ + H
₂O → H
₃O⁺ + SO²⁻
₄   K_a2 = 1.0×10⁻² 

HSO⁻
₄ is the bisulfate anion and SO²⁻
₄ is the sulfate anion. K_a1 and K_a2 are the acid dissociation constants. 

Because the hydration of sulfuric acid is thermodynamically favorable and the affinity of it for water is sufficiently strong, sulfuric acid is an excellent dehydrating agent. Concentrated sulfuric acid has a very powerful dehydrating property, removing water (H₂O) from other chemical compounds including sugar and other carbohydrates and producing carbon, heat, steam. 

In the laboratory, this is often demonstrated by mixing table sugar (sucrose) into sulfuric acid. The sugar changes from white to dark brown and then to black as carbon is formed. A rigid column of black, porous carbon will emerge as well. The carbon will smell strongly of caramel due to the heat generated.

${\displaystyle \overbrace {{\ce {C12H22O11}}} ^{\text{sucrose}}\ {\ce {->[{\ce {H2SO4}}]}}\ {\underset {\text{(black graphitic foam)}}{{\ce {12C}}}}+{\ce {11H2O}}_{\text{(g,l)}}}$

Similarly, mixing starch into concentrated sulfuric acid will give elemental carbon and water as absorbed by the sulfuric acid (which becomes slightly diluted). The effect of this can be seen when concentrated sulfuric acid is spilled on paper which is composed of cellulose; the cellulose reacts to give a burnt appearance, the carbon appears much as soot would in a fire. Although less dramatic, the action of the acid on cotton, even in diluted form, will destroy the fabric.

${\displaystyle \overbrace {\ce {(C6H10O5)_{\mathit {n}}}} ^{\text{polysaccharide}}\ {\ce {->[{\ce {H2SO4}}]}}\ 6n{\ce {C}}+5n{\ce {H2O}}}$

The reaction with copper(II) sulfate can also demonstrate the dehydration property of sulfuric acid. The blue crystal is changed into white powder as water is removed.

${\displaystyle \overbrace {\underset {\text{(blue crystal)}}{{\ce {CuSO4.5H2O}}}} ^{\text{copper(II) sulfate hydrate}}\ {\ce {->[{\ce {H2SO4}}]}}\ \overbrace {\underset {\text{(white powder)}}{{\ce {CuSO4}}}} ^{\text{Anhydrous copper(II) sulfate}}+{\ce {5H2O}}}$

Acid-base properties

As an acid, sulfuric acid reacts with most bases to give the corresponding sulfate. For example, the blue copper salt copper(II) sulfate, commonly used for electroplating and as a fungicide, is prepared by the reaction of copper(II) oxide with sulfuric acid:

CuO (s) + H
₂SO
₄ (aq) → CuSO
₄ (aq) + H
₂O (l)

Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate, for example, displaces acetic acid, CH
₃COOH, and forms sodium bisulfate:

H
₂SO
₄ + CH
₃COONa → NaHSO
₄ + CH
₃COOH

Similarly, reacting sulfuric acid with potassium nitrate can be used to produce nitric acid and a precipitate of potassium bisulfate. When combined with nitric acid, sulfuric acid acts both as an acid and a dehydrating agent, forming the nitronium ion NO⁺
₂, which is important in nitration reactions involving electrophilic aromatic substitution. This type of reaction, where protonation occurs on an oxygen atom, is important in many organic chemistry reactions, such as Fischer esterification and dehydration of alcohols. 

Solid state structure of the [D₃SO₄]⁺ ion present in [D₃SO₄]⁺[SbF₆]⁻, synthesized by using DF in place of HF. (see text)

 

When allowed to react with superacids, sulfuric acid can act as a base and be protonated, forming the [H₃SO₄]⁺ ion. Salt of [H₃SO₄]⁺ have been prepared using the following reaction in liquid HF:

((CH₃)₃SiO)₂SO₂ + 3 HF + SbF₅ → [H₃SO₄]⁺[SbF₆]⁻ + 2 (CH₃)₃SiF

The above reaction is thermodynamically favored due to the high bond enthalpy of the Si–F bond in the side product. Protonation using simply HF/SbF₅, however, have met with failure, as pure sulfuric acid undergoes self-ionization to give [H₃O]⁺ ions, which prevents the conversion of H₂SO₄ to [H₃SO₄]⁺ by the HF/SbF₅ system:

2 H₂SO₄ ⇌ [H₃O]⁺ + [HS₂O₇]⁻

Reactions with metals and strong oxidizing property

Dilute sulfuric acid reacts with metals via a single displacement reaction as with other typical acids, producing hydrogen gas and salts (the metal sulfate). It attacks reactive metals (metals at positions above copper in the reactivity series) such as iron, aluminium, zinc, manganese, magnesium, and nickel.

Fe (s) + H
₂SO
₄ (aq) → H
₂ (g) + FeSO
₄ (aq)

However, concentrated sulfuric acid is a strong oxidizing agent and does not react with metals in the same way as other typical acids. Sulfur dioxide, water and SO²⁻
₄ ions are evolved instead of the hydrogen and salts.

2 H₂SO₄ + 2 e⁻ → SO₂ + 2 H₂O + SO²⁻
₄

It can oxidize non-active metals such as tin and copper, depending upon the temperature.

Cu + 2 H₂SO₄ → SO₂ + 2 H₂O + SO²⁻
₄ + Cu²⁺

Lead and tungsten, however, are resistant to sulfuric acid.

Reactions with non-metals

Hot concentrated sulfuric acid oxidizes non-metals such as carbon (as bituminous coal) and sulfur.

C + 2 H₂SO₄ → CO₂ + 2 SO₂ + 2 H₂O

S + 2 H₂SO₄ → 3 SO₂ + 2 H₂O

Reaction with sodium chloride

It reacts with sodium chloride, and gives hydrogen chloride gas and sodium bisulfate:

NaCl + H₂SO₄ → NaHSO₄ + HCl

Electrophilic aromatic substitution

Benzene undergoes electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acids:

Occurrence

Rio Tinto with its highly acidic water

 

Pure sulfuric acid is not encountered naturally on Earth in anhydrous form, due to its great affinity for water. Dilute sulfuric acid is a constituent of acid rain, which is formed by atmospheric oxidation of sulfur dioxide in the presence of water – i.e., oxidation of sulfurous acid. Sulfur dioxide is the main byproduct produced when sulfur-containing fuels such as coal or oil are burned.

Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as iron sulfide. The resulting water can be highly acidic and is called acid mine drainage (AMD) or acid rock drainage (ARD). This acidic water is capable of dissolving metals present in sulfide ores, which results in brightly colored, toxic streams. The oxidation of pyrite (iron sulfide) by molecular oxygen produces iron(II), or Fe²⁺:

2 FeS
₂ (s) + 7 O
₂ + 2  H₂O → 2 Fe²⁺ + 4 SO²⁻
₄ + 4 H⁺

The Fe²⁺ can be further oxidized to Fe³⁺:

4 Fe²⁺ + O
₂ + 4 H⁺ → 4 Fe³⁺ + 2  H₂O

The Fe³⁺ produced can be precipitated as the hydroxide or hydrous iron oxide:

Fe³⁺ + 3  H₂O → Fe(OH)
₃↓ + 3 H⁺

The iron(III) ion ("ferric iron") can also oxidize pyrite:

FeS
₂(s) + 14 Fe³⁺ + 8  H₂O → 15 Fe²⁺ + 2 SO²⁻
₄ + 16 H⁺

When iron(III) oxidation of pyrite occurs, the process can become rapid. pH values below zero have been measured in ARD produced by this process. 

ARD can also produce sulfuric acid at a slower rate, so that the acid neutralizing capacity (ANC) of the aquifer can neutralize the produced acid. In such cases, the total dissolved solids (TDS) concentration of the water can be increased from the dissolution of minerals from the acid-neutralization reaction with the minerals. 

Sulfuric acid is used as a defense by certain marine species, for example, the phaeophyte alga Desmarestia munda (order Desmarestiales) concentrates sulfuric acid in cell vacuoles.

Stratospheric aerosol

In the stratosphere, the atmosphere's second layer that is generally between 10 and 50 km above Earth's surface, sulfuric acid is formed by the oxidation of volcanic sulfur dioxide by the hydroxyl radical:

SO
₂ + HO^• → HSO
₃

HSO
₃ + O
₂ → SO
₃ + HO
₂

SO
₃ +  H₂O → H
₂SO
₄

Because sulfuric acid reaches supersaturation in the stratosphere, it can nucleate aerosol particles and provide a surface for aerosol growth via condensation and coagulation with other water-sulfuric acid aerosols. This results in the stratospheric aerosol layer.

Extraterrestrial sulfuric acid

Venus

Sulfuric acid is produced in the upper atmosphere of Venus by the Sun's photochemical action on carbon dioxide, sulfur dioxide, and water vapor. Ultraviolet photons of wavelengths less than 169 nm can photodissociate carbon dioxide into carbon monoxide and atomic oxygen. Atomic oxygen is highly reactive. When it reacts with sulfur dioxide, a trace component of the Venusian atmosphere, the result is sulfur trioxide, which can combine with water vapor, another trace component of Venus's atmosphere, to yield sulfuric acid. In the upper, cooler portions of Venus's atmosphere, sulfuric acid exists as a liquid, and thick sulfuric acid clouds completely obscure the planet's surface when viewed from above. The main cloud layer extends from 45–70 km above the planet's surface, with thinner hazes extending as low as 30 km and as high as 90 km above the surface. The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain.

The atmosphere exhibits a sulfuric acid cycle. As sulfuric acid rain droplets fall down through the hotter layers of the atmosphere's temperature gradient, they are heated up and release water vapor, becoming more and more concentrated. When they reach temperatures above 300 °C, sulfuric acid begins to decompose into sulfur trioxide and water, both in the gas phase. Sulfur trioxide is highly reactive and dissociates into sulfur dioxide and atomic oxygen, which oxidizes traces of carbon monoxide to form carbon dioxide. Sulfur dioxide and water vapor rise on convection currents from the mid-level atmospheric layers to higher altitudes, where they will be transformed again into sulfuric acid, and the cycle repeats.

Europa

Infrared spectra taken by NASA's Galileo spacecraft show distinct absorptions on Jupiter's moon Europa that have been attributed to one or more sulfuric acid hydrates. Sulfuric acid in solution with water causes significant freezing-point depression of water's melting point, down to 210 K (−63 °C), and this would make the existence of liquid solutions beneath Europa's icy crust more likely. The interpretation of the spectra is somewhat controversial. Some planetary scientists prefer to assign the spectral features to the sulfate ion, perhaps as part of one or more minerals on Europa's surface.

Manufacture

Sulfuric acid is produced from sulfur, oxygen and water via the conventional contact process (DCDA) or the wet sulfuric acid process (WSA).

Contact process

In the first step, sulfur is burned to produce sulfur dioxide.

S (s) + O
₂ (g) → SO
₂ (g)

This is then oxidized to sulfur trioxide using oxygen in the presence of a vanadium(V) oxide catalyst. This reaction is reversible and the formation of the sulfur trioxide is exothermic.

2 SO
₂ (g) + O
₂ (g) ⇌ 2 SO
₃ (g) (in presence of V
₂O
₅)

The sulfur trioxide is absorbed into 97–98% H
₂SO
₄ to form oleum (H
₂S
₂O
₇), also known as fuming sulfuric acid. The oleum is then diluted with water to form concentrated sulfuric acid.

H
₂SO
₄ (l) + SO
₃ (g)→ H
₂S
₂O
₇ (l)

H
₂S
₂O
₇ (l) + H
₂O (l) → 2 H
₂SO
₄ (l)

Directly dissolving SO
₃ in water is not practical due to the highly exothermic nature of the reaction between sulfur trioxide and water. The reaction forms a corrosive aerosol that is very difficult to separate, instead of a liquid.

SO
₃ (g) + H
₂O (l) → H
₂SO
₄ (l)

Wet sulfuric acid process

In the first step, sulfur is burned to produce sulfur dioxide:

S(s) + O
₂(g) → SO
₂(g)

or, alternatively, hydrogen sulfide (H
₂S) gas is incinerated to SO
₂ gas:

2 H
₂S + 3 O
₂ → 2 H
₂O + 2 SO
₂ (−518 kJ/mol)

This is then oxidized to sulfur trioxide using oxygen with vanadium(V) oxide as catalyst.

2 SO
₂ + O
₂ → 2 SO
₃ (−99 kJ/mol) (reaction is reversible)

The sulfur trioxide is hydrated into sulfuric acid H
₂SO
₄:

SO
₃ + H
₂O → H
₂SO
₄(g) (−101 kJ/mol)

The last step is the condensation of the sulfuric acid to liquid 97–98% H
₂SO
₄:

H
₂SO
₄(g) → H
₂SO
₄(l) (−69 kJ/mol)

Other methods

Another method is the less well-known metabisulfite method, in which metabisulfite is placed at the bottom of a beaker, and 12.6 molar concentration hydrochloric acid is added. The resulting gas is bubbled through nitric acid, which will release brown/red vapors. The completion of the reaction is indicated by the ceasing of the fumes. This method does not produce an inseparable mist, which is quite convenient.

SO₂ + HNO₃ + H₂O→ H₂SO₄ + NO

Sulfuric acid can be produced in the laboratory by burning sulfur in air and dissolving the gas produced in a hydrogen peroxide solution.

SO₂ + H₂O₂ → H₂SO₄

Prior to 1900, most sulfuric acid was manufactured by the lead chamber process. As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants.

In early to mid nineteenth century "vitriol" plants existed, among other places, in Prestonpans in Scotland, Shropshire and the Lagan Valley in County Antrim Ireland where it was used as a bleach for linen. Early bleaching of linen was done using lactic acid from sour milk but this was a slow process and the use of vitriol sped up the bleaching process.

Uses

Sulfuric acid production in 2000

 

Sulfuric acid is a very important commodity chemical, and indeed, a nation's sulfuric acid production is a good indicator of its industrial strength. World production in 2004 was about 180 million tonnes, with the following geographic distribution: Asia 35%, North America (including Mexico) 24%, Africa 11%, Western Europe 10%, Eastern Europe and Russia 10%, Australia and Oceania 7%, South America 7%. Most of this amount (≈60%) is consumed for fertilizers, particularly superphosphates, ammonium phosphate and ammonium sulfates. About 20% is used in chemical industry for production of detergents, synthetic resins, dyestuffs, pharmaceuticals, petroleum catalysts, insecticides and antifreeze, as well as in various processes such as oil well acidifing, aluminum reduction, paper sizing, water treatment. About 6% of uses are related to pigments and include paints, enamels, printing inks, coated fabrics and paper, and the rest is dispersed into a multitude of applications such as production of explosives, cellophane, acetate and viscose textiles, lubricants, non-ferrous metals, and batteries.

Industrial production of chemicals

The major use for sulfuric acid is in the "wet method" for the production of phosphoric acid, used for manufacture of phosphate fertilizers. In this method, phosphate rock is used, and more than 100 million tonnes are processed annually. This raw material is shown below as fluorapatite, though the exact composition may vary. This is treated with 93% sulfuric acid to produce calcium sulfate, hydrogen fluoride (HF) and phosphoric acid. The HF is removed as hydrofluoric acid. The overall process can be represented as:

${\displaystyle {\ce {{\overset {fluorapatite}{Ca5F(PO4)3}}+{5H2SO4}+10H2O->{\overset {calcium~sulfate}{5CaSO4.2H2O}}+{HF}+3H3PO4}}}$

Ammonium sulfate, an important nitrogen fertilizer, is most commonly produced as a byproduct from coking plants supplying the iron and steel making plants. Reacting the ammonia produced in the thermal decomposition of coal with waste sulfuric acid allows the ammonia to be crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry. 

Another important use for sulfuric acid is for the manufacture of aluminium sulfate, also known as paper maker's alum. This can react with small amounts of soap on paper pulp fibers to give gelatinous aluminium carboxylates, which help to coagulate the pulp fibers into a hard paper surface. It is also used for making aluminium hydroxide, which is used at water treatment plants to filter out impurities, as well as to improve the taste of the water. Aluminium sulfate is made by reacting bauxite with sulfuric acid:

2 AlO(OH) + 3 H
₂SO
₄ → Al
₂(SO
₄)
₃ + 4 H
₂O

Sulfuric acid is also important in the manufacture of dyestuffs solutions.

Sulfur–iodine cycle

The sulfur–iodine cycle is a series of thermo-chemical processes used to obtain hydrogen. It consists of three chemical reactions whose net reactant is water and whose net products are hydrogen and oxygen. Step one of cycle is the Bunsen reaction.

2 H 2SO 4 → 2 SO 2 + 2 H 2O + O 2		(830 °C)
I 2 + SO 2 + 2 H 2O → 2 HI + H 2SO 4		(120 °C)
2 HI → I 2 + H 2		(320 °C)

The sulfur and iodine compounds are recovered and reused, hence the consideration of the process as a cycle. This process is endothermic and must occur at high temperatures, so energy in the form of heat has to be supplied. 

The sulfur–iodine cycle has been proposed as a way to supply hydrogen for a hydrogen-based economy. It does not require hydrocarbons like current methods of steam reforming. But note that all of the available energy in the hydrogen so produced is supplied by the heat used to make it.

The sulfur–iodine cycle is currently being researched as a feasible method of obtaining hydrogen, but the concentrated, corrosive acid at high temperatures poses currently insurmountable safety hazards if the process were built on a large scale.

Industrial cleaning agent

Sulfuric acid is used in large quantities by the iron and steelmaking industry to remove oxidation, rust, and scaling from rolled sheet and billets prior to sale to the automobile and major appliances industry. Used acid is often recycled using a spent acid regeneration (SAR) plant. These plants combust spent acid with natural gas, refinery gas, fuel oil or other fuel sources. This combustion process produces gaseous sulfur dioxide (SO
₂) and sulfur trioxide (SO
₃) which are then used to manufacture "new" sulfuric acid. SAR plants are common additions to metal smelting plants, oil refineries, and other industries where sulfuric acid is consumed in bulk, as operating a SAR plant is much cheaper than the recurring costs of spent acid disposal and new acid purchases. 

Hydrogen peroxide (H
₂O
₂) can be added to sulfuric acid to produce piranha solution, a powerful but very toxic cleaning solution with which substrate surfaces can be cleaned. Piranha solution is typically used in the microelectronics industry, and also in laboratory settings to clean glassware.

Catalyst

Sulfuric acid is used for a variety of other purposes in the chemical industry. For example, it is the usual acid catalyst for the conversion of cyclohexanone oxime to caprolactam, used for making nylon. It is used for making hydrochloric acid from salt via the Mannheim process. Much H
₂SO
₄ is used in petroleum refining, for example as a catalyst for the reaction of isobutane with isobutylene to give isooctane, a compound that raises the octane rating of gasoline (petrol). Sulfuric acid is also often used as a dehydrating or oxidizing agent in industrial reactions, such as the dehydration of various sugars to form solid carbon.

Electrolyte

Acidic drain cleaners usually contain sulfuric acid at a high concentration which turns a piece of pH paper red and chars it instantly, demonstrating both the strong acidic nature and dehydrating property.

 

Sulfuric acid acts as the electrolyte in lead–acid batteries (lead-acid accumulator):

At anode:

Pb + SO
₄²⁻ ⇌ PbSO
₄ + 2 e⁻

At cathode:

PbO
₂ + 4 H⁺ + SO
₄²⁻ + 2 e⁻ ⇌ PbSO
₄ + 2 H₂O

An acidic drain cleaner can be used to dissolve grease, hair and even tissue paper inside water pipes.

 

Overall:

Pb + PbO
₂ + 4 H⁺ + 2 SO
₄²⁻ ⇌ 2 PbSO
₄ + 2 H₂O

Domestic uses

Sulfuric acid at high concentrations is frequently the major ingredient in acidic drain cleaners which are used to remove grease, hair, tissue paper, etc. Similar to their alkaline versions, such drain openers can dissolve fats and proteins via hydrolysis. Moreover, as concentrated sulfuric acid has a strong dehydrating property, it can remove tissue paper via dehydrating process as well. Since the acid may react with water vigorously, such acidic drain openers should be added slowly into the pipe to be cleaned.

History

John Dalton's 1808 sulfuric acid molecule shows a central sulfur atom bonded to three oxygen atoms, or sulfur trioxide, the anhydride of sulfuric acid.

 

The study of vitriol, a category of glassy minerals from which the acid can be derived, began in ancient times. Sumerians had a list of types of vitriol that they classified according to the substances' color. Some of the earliest discussions on the origin and properties of vitriol is in the works of the Greek physician Dioscorides (first century AD) and the Roman naturalist Pliny the Elder (23–79 AD). Galen also discussed its medical use. Metallurgical uses for vitriolic substances were recorded in the Hellenistic alchemical works of Zosimos of Panopolis, in the treatise Phisica et Mystica, and the Leyden papyrus X.

Medieval Islamic era alchemists, Jābir ibn Hayyān (c. 721 – c. 815 AD, also known as Geber), Razi (865 – 925 AD), and Jamal Din al-Watwat (d. 1318, wrote the book Mabāhij al-fikar wa-manāhij al-'ibar), included vitriol in their mineral classification lists. Ibn Sina focused on its medical uses and different varieties of vitriol.

Sulfuric acid was called "oil of vitriol" by medieval European alchemists because it was prepared by roasting "green vitriol" (iron(II) sulfate) in an iron retort. There are references to it in the works of Vincent of Beauvais and in the Compositum de Compositis ascribed to Saint Albertus Magnus. A passage from Pseudo-Geber´s Summa Perfectionis was long considered to be the first recipe for sulfuric acid, but this was a misinterpretation.

In the seventeenth century, the German-Dutch chemist Johann Glauber prepared sulfuric acid by burning sulfur together with saltpeter (potassium nitrate, KNO
₃), in the presence of steam. As saltpeter decomposes, it oxidizes the sulfur to SO
₃, which combines with water to produce sulfuric acid. In 1736, Joshua Ward, a London pharmacist, used this method to begin the first large-scale production of sulfuric acid. 

In 1746 in Birmingham, John Roebuck adapted this method to produce sulfuric acid in lead-lined chambers, which were stronger, less expensive, and could be made larger than the previously used glass containers. This process allowed the effective industrialization of sulfuric acid production. After several refinements, this method, called the lead chamber process or "chamber process", remained the standard for sulfuric acid production for almost two centuries.

Sulfuric acid created by John Roebuck's process approached a 65% concentration. Later refinements to the lead chamber process by French chemist Joseph Louis Gay-Lussac and British chemist John Glover improved concentration to 78%. However, the manufacture of some dyes and other chemical processes require a more concentrated product. Throughout the 18th century, this could only be made by dry distilling minerals in a technique similar to the original alchemical processes. Pyrite (iron disulfide, FeS
₂) was heated in air to yield iron(II) sulfate, FeSO
₄, which was oxidized by further heating in air to form iron(III) sulfate, Fe₂(SO₄)₃, which, when heated to 480 °C, decomposed to iron(III) oxide and sulfur trioxide, which could be passed through water to yield sulfuric acid in any concentration. However, the expense of this process prevented the large-scale use of concentrated sulfuric acid.

In 1831, British vinegar merchant Peregrine Phillips patented the contact process, which was a far more economical process for producing sulfur trioxide and concentrated sulfuric acid. Today, nearly all of the world's sulfuric acid is produced using this method.

Safety

Laboratory hazards

Drops of 98% sulfuric acid char a piece of tissue paper instantly. Carbon is left after the dehydration reaction staining the paper black.

 

Superficial chemical burn caused by two 98% sulfuric acid splashes (forearm skin)

 

Sulfuric acid is capable of causing very severe burns, especially when it is at high concentrations. In common with other corrosive acids and alkali, it readily decomposes proteins and lipids through amide and ester hydrolysis upon contact with living tissues, such as skin and flesh. In addition, it exhibits a strong dehydrating property on carbohydrates, liberating extra heat and causing secondary thermal burns. Accordingly, it rapidly attacks the cornea and can induce permanent blindness if splashed onto eyes. If ingested, it damages internal organs irreversibly and may even be fatal.^[5] Protective equipment should hence always be used when handling it. Moreover, its strong oxidizing property makes it highly corrosive to many metals and may extend its destruction on other materials. Because of such reasons, damage posed by sulfuric acid is potentially more severe than that by other comparable strong acids, such as hydrochloric acid and nitric acid.

Sulfuric acid must be stored carefully in containers made of nonreactive material (such as glass). Solutions equal to or stronger than 1.5 M are labeled "CORROSIVE", while solutions greater than 0.5 M but less than 1.5 M are labeled "IRRITANT". However, even the normal laboratory "dilute" grade (approximately 1 M, 10%) will char paper if left in contact for a sufficient time.

The standard first aid treatment for acid spills on the skin is, as for other corrosive agents, irrigation with large quantities of water. Washing is continued for at least ten to fifteen minutes to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing is removed immediately and the underlying skin washed thoroughly.

Dilution hazards

Preparation of the diluted acid can be dangerous due to the heat released in the dilution process. To avoid splattering, the concentrated acid is usually added to water and not the other way around. Water has a higher heat capacity than the acid, and so a vessel of cold water will absorb heat as acid is added. 

Comparison of sulfuric acid and water

Physical property	H2SO4	Water	Units
Density	1.84	1.0	kg/L
Volumetric heat capacity	2.54	4.18	kJ/L
Boiling point	337	100	°C

Also, because the acid is denser than water, it sinks to the bottom. Heat is generated at the interface between acid and water, which is at the bottom of the vessel. Acid will not boil, because of its higher boiling point. Warm water near the interface rises due to convection, which cools the interface, and prevents boiling of either acid or water. 

In contrast, addition of water to concentrated sulfuric acid results in a thin layer of water on top of the acid. Heat generated in this thin layer of water can boil, leading to the dispersal of a sulfuric acid aerosol or worse, an explosion. 

Preparation of solutions greater than 6 M (35%) in concentration is most dangerous, because the heat produced may be sufficient to boil the diluted acid: efficient mechanical stirring and external cooling (such as an ice bath) are essential. 

Reaction rates double for about every 10 degree Celsius increase in temperature. Therefore, the reaction will become more violent as dilution proceeds, unless the mixture is given time to cool. Adding acid to warm water will cause a violent reaction. 

On a laboratory scale, sulfuric acid can be diluted by pouring concentrated acid onto crushed ice made from de-ionized water. The ice melts in an endothermic process while dissolving the acid. The amount of heat needed to melt the ice in this process is greater than the amount of heat evolved by dissolving the acid so the solution remains cold. After all the ice has melted, further dilution can take place using water.

Industrial hazards

Sulfuric acid is non-flammable. 

The main occupational risks posed by this acid are skin contact leading to burns (see above) and the inhalation of aerosols. Exposure to aerosols at high concentrations leads to immediate and severe irritation of the eyes, respiratory tract and mucous membranes: this ceases rapidly after exposure, although there is a risk of subsequent pulmonary edema if tissue damage has been more severe. At lower concentrations, the most commonly reported symptom of chronic exposure to sulfuric acid aerosols is erosion of the teeth, found in virtually all studies: indications of possible chronic damage to the respiratory tract are inconclusive as of 1997. Repeated occupational exposure to sulfuric acid mists may increase the chance of lung cancer by up to 64 percent. In the United States, the permissible exposure limit (PEL) for sulfuric acid is fixed at 1 mg/m³: limits in other countries are similar. There have been reports of sulfuric acid ingestion leading to vitamin B12 deficiency with subacute combined degeneration. The spinal cord is most often affected in such cases, but the optic nerves may show demyelination, loss of axons and gliosis.

Legal restrictions

International commerce of sulfuric acid is controlled under the United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances, 1988, which lists sulfuric acid under Table II of the convention as a chemical frequently used in the illicit manufacture of narcotic drugs or psychotropic substances.

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.