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Sulfuric Acid (H2SO4)

Sulfuric acid


Systematic name

Sulfuric Acid

Other names

Battery acid


Oil of vitriol

Molecular formula

H2SO4 (aq)

Molar mass

98.08 g mol−1


Clear, colorless,
odorless liquid

CAS number



Density and phase

1.84 g cm−3, liquid

Solubility in water

Fully miscible

Melting point

10 °C (283 K)

Boiling point

338 °C (611 K)






26.7 cP at 20°C




EU classification

Corrosive (C)

NFPA 704

R (risk)-phrases

R35 (causes severe burns)


S1/2, S26, S30, S45

Flash point


RTECS number


Supplementary data

Structure & properties

n, εr, etc.

Thermodynamic data

Phase behavior
Solid, liquid, gas

Spectral data


Related compounds

Related strong acids

Selenic acid
Hydrochloric acid
Nitric acid

Related compounds

Hydrogen sulfide
Sulfurous acid

Peroxymonosulfuric acid

Sulfur trioxide

GB Disclaimer: Except where noted otherwise, data is given for materials in their standard state (at 25 °C, 100 kPa)

Sulfuric acid, H2SO4, is a strong mineral acid. It is soluble in water at all concentrations. It was once known as oil of vitriol, and has an oily consistency when concentrated.

Physical properties:

Forms of Sulfuric acid

Although nearly 100% sulfuric acid can be made, this loses SO3 at the boiling point to produce 98.3% acid. The 98% 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:

·        10%: Dilute sulfuric acid for laboratory use

·        33.5%: Battery acid (used in automotive batteries)

·        37.52%: Battery acid (used in traction, lift truck, forklift batteries). 1.285 specific gravity (spgr).

·        62.18%: Chamber or fertilizer acid

·        77.67%: Tower or Glover acid

·        98%: Concentrated acid

Since sulfuric acid is a strong acid, a 0.50 M solution of sulfuric acid has a pH close to zero.



Industrial hazards

Although sulfuric acid is non-flammable, contact with metals in the event of a spillage can lead to the liberation of hydrogen gas. The dispersal of acid aerosols and gaseous sulfur dioxide is an additional hazard of fires involving sulfuric acid. Water should not be used as the extinguishing agent because of the risk of further dispersal of aerosols: carbon dioxide is preferred where possible.

Sulfuric acid is not considered toxic besides its obvious corrosive hazard, and the main occupational risks are skin contact leading to burns (see below) 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. In the United States, the permissible exposure limit (PEL) for sulfuric acid is fixed at 1 mg/m3: limits in other countries are similar. Interestingly 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.


Laboratory hazards

The corrosive properties of sulfuric acid are accentuated by its highly exothermic reaction with water. Hence burns from sulfuric acid are potentially more serious than those of comparable strong acids (e.g. hydrochloric acid), as there is additional tissue damage due to dehydration and particularly due to the heat liberated by the reaction with water, i.e. secondary thermal damage. The danger is obviously greater with more concentrated preparations of sulfuric acid, but it should be remembered that even the normal laboratory "dilute" grade (approx. 1 M, 10%) will char paper by dehydration if left in contact for a sufficient length of time. Solutions equal to or stronger than 1.5 M should be labeled CORROSIVE, while solutions 0.5 M but less than 1.5 M should be labeled IRRITANT. The standard first aid treatment for acid spills on the skin is, as for other corrosive agents, irrigation with large quantities of water: Washing should be continued for a sufficient length of time—at least ten to fifteen minutes—in order to cool the tissue surrounding the acid burn and to prevent secondary damage. Contaminated clothing must be removed immediately and the underlying skin washed thoroughly.

Preparation of the diluted acid can also be dangerous due to the heat released in the dilution process. It is essential that the concentrated acid is added to water and not the other way round, to take advantage of the relatively high heat capacity of water. Addition of water to concentrated sulfuric acid leads at best to the dispersal of a sulfuric acid aerosol, at worst to an explosion. Preparation of solutions greater than 6 M (35%) in concentration is the most dangerous, as the heat produced can be sufficient to boil the diluted acid: efficient mechanical stirring and external cooling (e.g. an ice bath) are essential.


Potential Health Hazards

SKIN: Causes severe burns.

EYES: Liquid contact can cause irritation, corneal burns, and conjunctivitis. May result in severe or permanent injury. May cause blindness.

INHALATION: Inhalation of fumes or acid mist can cause irritation or corrosive burns to the upper respiratory system, including the nose, mouth and throat. May irritate the lungs. May cause pulmonary edema.

INGESTION: Causes burns of the mouth, throat and stomach. May be fatal if swallowed. Hazards are also applicable to dilute solutions.


Polarity and conductivity

Anhydrous H2SO4 is a very polar liquid, with a dielectric constant of around 100. This is due to the fact that it can dissociate by protonating itself, a process known as autoprotolysis, which occurs to a high degree, more than 10 billion times the level seen in water:

2 H2SO4 H3SO4+ + HSO4


This allows protons to be highly mobile in H2SO4. It also makes sulfuric acid excellent for many reactions. In fact, the equilibrium is more complex than shown above. 100% H2SO4 contains the following species at equilibrium (figures shown as mol per kg solvent):

HSO4 (15.0), H3SO4+ (11.3), H3O+ (8.0), HS2O7 (4.4), H2S2O7 (3.6), H2O (0.1)



A mixture of sulfuric acid and water is used as the electrolyte in lead-acid battery where it undergoes a reversible reaction where lead and lead dioxide are converted to lead(II) sulfate.

Besides it’s use in batteries, sulfuric acid is a very important commodity chemical. A nation's sulfuric acid production is a good indicator of its industrial strength. The major use (60% of total worldwide) for sulfuric acid is in the "wet method" for the production of phosphoric acid, used for manufacture of phosphate fertilizers as well as trisodium phosphate for detergents. In this method phosphate rock is used, and more than 100 million tonnes is 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:

Ca5F(PO4)3 + 5 H2SO4 + 10 H2O 5 CaSO42 H2O + HF + 3 H3PO4

Sulfuric acid is used in large quantities in iron and steel making principally as pickling-acid used to remove oxidation, rust and scale from rolled sheet and billets prior to sale into the automobile and white-goods business. The used acid is often re-cycled using a Spent Acid Regeneration (SAR) plant. These plants combust the spent acid with natural gas, refinery gas, fuel oil or other suitable fuel source. This combustion process produces gaseous sulfur dioxide (SO2) and sulfur trioxide (SO3) which are then used to manufacture "new" sulfuric acid. These types of plants are common additions to metal smelting plants, oil refineries, and other places where sulfuric acid is consumed on a large scale, as operating a SAR plant is much cheaper than purchasing the commodity on the open market.

Ammonium sulfate, an important nitrogen fertilizer is most commonly produced as a by-product 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 crystallised 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 papermaker's alum. This can react with small amounts of soap on paper pulp fibres to give gelatinous aluminium carboxylates, which help to coagulate the pulp fibres 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:

Al2O3 + 3 H2SO4 Al2(SO4)3 + 3 H2O

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 cyclohexanoneoxime to caprolactam, used for making nylon. It is used for making hydrochloric acid from salt via the Mannheim process. Much H2SO4 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 important in the manufacture of dyestuffs.

Sulfuric acid is also the principal ingredient in some drain cleaners, used to clear blockages consisting of paper, rags, and other materials not easily dissolved by caustic solutions.

Sulfuric acid is also used as a general dehydrating agent in its concentrated form. See Reaction with water.


Chemical properties

Reaction with water

The hydration reaction of sulfuric acid is highly exothermic. If water is added to concentrated sulfuric acid, it can boil and spit dangerously. One should always add the acid to the water rather than the water to the acid. This can be remembered through mnemonics such as "Always do things as you oughta, add the acid to the water. If you think your life's too placid, add the water to the acid", "A.A.: Add Acid", or "Drop acid, not water." Note that part of this problem is due to the relative densities of the two liquids. Water is less dense than sulfuric acid and will tend to float above the acid. The reaction is best thought of as forming hydronium ions, by:

H2SO4 + H2O H3O+ + HSO4-

And then:

HSO4- + H2O H3O+ + SO42-

Other reactions of sulfuric acid

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

CuO + H2SO4 CuSO4 + H2O


Sulfuric acid can be used to displace weaker acids from their salts, for example sodium acetate gives acetic acid:


Likewise the reaction of sulfuric acid with potassium nitrate can be used to produce nitric acid, along with a precipitate of potassium bisulfate. With nitric acid itself, sulfuric acid acts as both an acid and a dehydrating agent, forming the nitronium ion NO2+, 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 reactions in organic chemistry, such as Fischer esterification and dehydration of alcohols.

Sulfuric acid reacts with most metals in a single displacement reaction to produce hydrogen gas and the metal sulfate. Dilute H2SO4 attacks iron, aluminium, zinc, manganese and nickel, but tin and copper require hot concentrated acid. Lead and tungsten are, however, resistant to sulfuric acid. The reaction with iron (shown) is typical for most of these metals, but the reaction with tin is unusual in that it produces sulfur dioxide rather than hydrogen.

Fe(s) + H2SO4(aq) H2(g) + FeSO4(aq)

Sn(s) + 2 H2SO4(aq) SnSO4(aq) + 2 H2O(l) + SO2(g)


Sulfuric acid is produced from sulfur, oxygen and water via the contact process.

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

(1) S(s) + O2(g) SO2(g)


This is then oxidised to sulfur trioxide using oxygen in the presence of a vanadium(V) oxide catalyst.

(2) 2 SO2 + O2(g) 2 SO3(g)     (in presence of V2O5)


Finally, the sulfur trioxide is treated with water (usually as 97-98% H2SO4 containing 2-3% water) to produce 98-99% sulfuric acid.

(3) SO3(g) + H2O(l) H2SO4(l)


Note that directly dissolving SO3 in water is impractical due to the highly exothermic nature of the reaction. Mists are formed instead of a liquid. Alternatively, the SO3 is absorbed into H2SO4 to produce oleum (H2S2O7), which is then diluted to form sulfuric acid.

(3) H2SO4(l) + SO3 H2S2O7(l)


S-Phrases / Sulfuric Acid

S1: Keep locked up.

S2: Keep out of the reach of children.

S26: In case of contact with eyes, rinse immediately with plenty of water and seek medical advice.

S30: Never add water to this product.

S45: In case of accident or if you feel unwell, seek medical advice immediately (show the label where possible).



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Last Updated: Monday, December 03, 2007 - 7:18 AM Eastern Time.