Sulfuric acid

Although nearly 100% sulfuric acid solutions 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:[12][13]

"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 (<70% to avoid contamination with nitrosylsulfuric acid) and tower acid being the acid recovered from the bottom of the Glover tower.[12][13] 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.[13]

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 <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C,[14] and 98% sulfuric acid has a <1 mmHg vapor pressure at 40 °C.[15]

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 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).

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.[16]

2 H
₂SO
₄ ⇌ H
₃SO⁺
₄ + HSO⁻
₄

The equilibrium constant for the autoprotolysis is[16]

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.

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.[17] 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 ≈ 10³ (strong acid)

HSO⁻
₄ + H
₂O → H
₃O⁺
+ SO²⁻
₄   K_a2 = 1.0×10⁻² [18]

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, and 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.[19]

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

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.

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

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:[20]

2 H₂SO₄ ⇌ [H₃O]⁺ + [HS₂O₇]⁻

Even dilute sulfuric acid reacts with many 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 + H
₂SO
₄ → H
₂ + FeSO
₄

Concentrated sulfuric acid can serve as an oxidizing agent, releasing sulfur dioxide:[6]

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

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

Hot concentrated sulfuric acid oxidizes carbon[21] (as bituminous coal) and sulfur.

C + 2 H₂SO₄ → CO₂ + 2 SO₂ + 2 H₂O

S + 2 H₂SO₄ → 3 SO₂ + 2 H₂O

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

NaCl + H₂SO₄ → NaHSO₄ + HCl

Benzene undergoes electrophilic aromatic substitution with sulfuric acid to give the corresponding sulfonic acids:[22]

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:[24]

SO
₂ + HO^• → HSO
₃

HSO
₃ + O
₂ → SO
₃ + HO
₂

SO
₃ + H
2O → 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.[24]

The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain.[25] Jupiter's moon Europa is also thought to have an atmosphere containing sulfuric acid hydrates.[26]

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

S (s) + O
₂ → SO
₂

The sulfur dioxide is oxidized to sulfur trioxide by 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 + O
₂ ⇌ 2 SO
₃

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 + SO
₃ → H
₂S
₂O
₇

H
₂S
₂O
₇ + H
₂O → 2 H
₂SO
₄

Directly dissolving SO
₃ in water is not practiced.

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

S + O
₂ → SO
₂

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)

The sulfur dioxide 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)

A method that is the less well-known is the 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 of nitrogen dioxide as the reaction proceeds. 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.

3SO₂ + 2HNO₃ + 2H₂O → 3H₂SO₄ + 2NO

In principle, sulfuric acid can be produced in the laboratory by burning sulfur in air followed by dissolving the resulting sulfur dioxide in a hydrogen peroxide solution.

SO₂ + H₂O₂ → H₂SO₄

Alternatively, dissolving sulfur dioxide in an aqueous solution of an oxidizing metal salt such as copper (II) or iron (III) chloride:

2FeCl₃ + 2H₂O + SO₂ → 2FeCl₂ + H₂SO₄ + 2HCl

2CuCl₂ + 2H₂O + SO₂ → 2CuCl + H₂SO₄ + 2HCl

Two less well-known laboratory methods of producing sulfuric acid, albeit in dilute form and requiring some extra effort in purification. A solution of copper (II) sulfate can be electrolyzed with a copper cathode and platinum/graphite anode to give spongy copper at cathode and evolution of oxygen gas at the anode, the solution of dilute sulfuric acid indicates completion of the reaction when it turns from blue to clear (production of hydrogen at cathode is another sign):

2CuSO₄ + 2H₂O → 2Cu + 2H₂SO₄ + O₂

More costly, dangerous, and troublesome yet novel is the electrobromine method, which employs a mixture of sulfur, water, and hydrobromic acid as the electrolytic solution. The sulfur is pushed to bottom of container under the acid solution, then the copper cathode and platinum/graphite anode are used with the cathode near the surface and the anode is positioned at bottom of the electrolyte to apply the current. This may take longer and emits toxic bromine/sulfur bromide vapors, but the reactant acid is recyclable, overall only the sulfur and water are converted to sulfuric acid (omitting losses of acid as vapors):

2HBr → H₂ + Br₂ (electrolysis of aqueous hydrogen bromide)

Br₂ + Br⁻ ↔ Br₃⁻ (initial tribromide production, eventually reverses as Br⁻ depletes)

2S + Br₂ → S₂Br₂ (bromine reacts with sulfur to form disulfur dibromide)

S₂Br₂ + 8H₂O + 5Br₂ → 2H₂SO₄ + 12HBr (oxidation and hydration of disulfur dibromide)

Prior to 1900, most sulfuric acid was manufactured by the lead chamber process.[27] As late as 1940, up to 50% of sulfuric acid manufactured in the United States was produced by chamber process plants.

In the 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.[28]

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.

The sulfur–iodine cycle is a series of thermo-chemical processes possibly usable to produce hydrogen from water. It consists of three chemical reactions whose net reactant is water and whose net products are hydrogen and oxygen.

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

The compounds of sulfur and iodine 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 is an alternative to electrolysis, and 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.[31][32]

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.

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 oxidising agent in industrial reactions, such as the dehydration of various sugars to form solid carbon.

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

Sulfuric acid at high concentrations is frequently the major ingredient in acidic drain cleaners[11] 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.

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.[6][7] 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.[6] 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.

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.

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.[37] 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.

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.[38] 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.