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Sulfuric acid

2007 Schools Wikipedia Selection. Related subjects: Chemical compounds

                            Sulfuric acid
                    Sulfuric acid Sulfuric acid
                               General
       Systematic name                 Sulfuric acid (Sulphuric acid)
           Other names                                 Oil of Vitriol
     Molecular formula                                 H[2]SO[4] (aq)
            Molar mass                                    98.08 g/mol
            Appearance                              Clear, colorless,
                                                      odorless liquid
            CAS number                                    [7664-93-9]
                             Properties
     Density and phase                            1.84 g/cm^3, liquid
   Solubility in water                                 Fully miscible
                                                         (exothermic)
         Melting point                                   10°C (283 K)
         Boiling point                                  337°C (610 K)
                 pK[a]                                         −3.0
                                                                 1.99
             Viscosity                               26.7 c P at 20°C
                               Hazards
                  MSDS                                  External MSDS
     EU classification                                 Corrosive (C),
                                                            Toxic (T)
              NFPA 704

                       0
                       3
                       2
                       [DEL: W :DEL]

             R-phrases                            R23, R24, R25, R35,
                                                   R36, R37, R38, R49
             S-phrases                            S23, S30, S36, S37,
                                                             S39, S45
           Flash point                                  Non-flammable
          RTECS number                                      WS5600000
                       Supplementary data page
Structure & properties                                  n, ε[r], etc.
    Thermodynamic data                                Phase behaviour
                                                   Solid, liquid, gas
         Spectral data                                UV, IR, NMR, MS
                          Related compounds
  Related strong acids              Selenic acid
                                       Hydrochloric acid
                                             Nitric acid
                                                      Phosphoric acid
     Related compounds             Hydrogen sulfide
                                           Sulfurous acid
                                   Peroxymonosulfuric acid
                                          Sulfur trioxide
                                                                Oleum
          Except where noted otherwise, data are given for
                materials in their standard state (at 25°C, 100 kPa)
                                    Infobox disclaimer and references

   Sulfuric acid (British English: sulphuric acid), H[2]SO[4], is a strong
   mineral acid. It is soluble in water at all concentrations. It was once
   known as oil of vitriol, coined by the 8th-century Alchemist Jabir ibn
   Hayyan, the chemical's probable discoverer. Sulfuric acid has many
   applications, and is produced in greater amounts than any other
   chemical besides water. World production in 2001 was 165 million
   tonnes, with an approximate value of $8 billion. Principal uses include
   ore processing, fertilizer manufacturing, oil refining, wastewater
   processing, and chemical synthesis. Many proteins are made of
   sulfur-containing amino acids (like cysteine and methionine) which
   produce sulfuric acid when metabolized by the body.

Physical properties

Forms of sulfuric acid

   Although 100% sulfuric acid can be made, this loses SO[3] at the
   boiling point to produce 98.3% acid. The 98% grade is also more stable
   for storage, making it the usual form for "concentrated" sulfuric acid.
   Other concentrations of sulfuric acid are used for different purposes.
   Some common concentrations are:
     * 10%, dilute sulfuric acid for laboratory use (pH 1)
     * 33.5%, battery acid (used in lead-acid batteries) (pH 0.5)
     * 62.18%, chamber or fertilizer acid (pH about 0.4)
     * 77.67%, tower or Glover acid (pH about 0.25)
     * 98%, concentrated (pH about 0.1)

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

   Different purities are also available. Technical grade H[2]SO[4] is
   impure and often colored, but it is suitable for making fertilizer.
   Pure grades such as US Pharmacopoeia (USP) grade are used for making
   pharmaceuticals and dyestuffs.

   When high concentrations of SO[3](g) are added to sulfuric acid,
   H[2]S[2]O[7] forms. This is called pyrosulfuric acid, fuming sulfuric
   acid or oleum or, less commonly, Nordhausen acid. Concentrations of
   oleum are either expressed in terms of % SO[3] (called % oleum) or as %
   H[2]SO[4] (the amount made if H[2]O were added); common concentrations
   are 40% oleum (109% H[2]SO[4]) and 65% oleum (114.6% H[2]SO[4]). Pure
   H[2]S[2]O[7] is in fact a solid, melting point 36 °C.

Polarity and conductivity

   Anhydrous H[2]SO[4] 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 H[2]SO[4] ⇌ H[3]SO[4]^+ + HSO[4]^−

   This allows protons to be highly mobile in H[2]SO[4]. It also makes
   sulfuric acid an excellent solvent for many reactions. In fact, the
   equilibrium is more complex than shown above. 100% H[2]SO[4] contains
   the following species at equilibrium (figures shown as mmol per kg
   solvent): HSO[4]^− (15.0), H[3]SO[4]^+ (11.3), H[3]O^+ (8.0),
   HS[2]O[7]^− (4.4), H[2]S[2]O[7] (3.6), H[2]O (0.1).

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:

          H[2]SO[4] + H[2]O → H[3]O^+ + HSO[4]^-

   And then:

          HSO[4]^- + H[2]O → H[3]O^+ + SO[4]^2-

   Because the hydration of sulfuric acid is thermodynamically favorable (
   ΔH = -880 k J/ mol), sulfuric acid is an excellent dehydrating agent,
   and is used to prepare many dried fruits. The affinity of sulfuric acid
   for water is sufficiently strong that it will take hydrogen and oxygen
   atoms out of other compounds; for example, mixing starch
   (C[6]H[12]O[6])[n] and concentrated sulfuric acid will give elemental
   carbon and water which is absorbed by the sulfuric acid (which becomes
   slightly diluted): (C[6]H[12]O[6])[n] → 6C + 6H[2]O. The effect of this
   can be seen when concentrated sulfuric acid is spilled on paper; the
   starch reacts to give a burned appearance, the carbon appears as soot
   would in a fire. A more dramatic illustration occurs when sulfuric acid
   is added to a tablespoon of white sugar in a cup when a tall rigid
   column of black porous carbon smelling strongly of caramel emerges from
   the cup.

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 + H[2]SO[4] → CuSO[4] + H[2]O

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

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

   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 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 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 H[2]SO[4] 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) + H[2]SO[4](aq) → H[2](g) + FeSO[4](aq)

          Sn(s) + 2 H[2]SO[4](l) → SnSO[4] + 2 H[2]O + SO[2]

Environmental aspects

   Sulfuric acid is a constituent of acid rain, being formed by
   atmospheric oxidation of sulfur dioxide in the presence of water - i.e.
   oxidation of sulfurous acid. Sulfur dioxide is the main product when
   the sulphur in sulfur-containing fuels such as coal or oil is 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 Rock Drainage (ARD). The acidic water so formed can
   dissolve metals present in sulfide ores, resulting in brightly colored
   and toxic streams. The oxidation of iron sulfide pyrite by molecular
   oxygen produces iron(II), or Fe^2+:

          FeS[2] + 7/2 O[2] + H[2]O → Fe^2+ + 2 SO[4]^2- + 2 H^+

   The Fe^2+ can be further oxidized to Fe^3+, according to:

          Fe^2+ + 1/4 O[2] + H^+ → Fe^3+ + 1/2 H[2]O

   and the Fe^3+ so produced can be precipitated as the hydroxide or
   hydrous oxide. The equation for the formation of the hydroxide is:

          Fe^3+ + 3 H[2]O → Fe(OH)[3] + 3 H^+

   The iron(III) ion ("ferric iron", in casual nomenclature)can also
   oxidize pyrite. When iron(III) oxidation of pyrite occurs, the process
   can become rapid and pH values below zero have been measured in ARD
   from this process.

   ARD can also produce sulfuric acid at a slower rate, so that the Acid
   Neutralization 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 form the dissolution of
   minerals from the acid-neutralization reaction with the minerals.

Extraterrestrial sulfuric acid

   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' atmosphere, to yield sulfuric acid.

          CO[2] → CO + O
          SO[2] + O → SO[3]
          SO[3] + H[2]O → H[2]SO[4]

   In the upper, cooler portions of Venus' atmosphere, sulfuric acid can
   exist as a liquid, and thick sulfuric acid clouds completely obscure
   the planet's surface from above. The main cloud layer extends from
   45–70 km above the planet's surface, with thinner hazes extending as
   low as 30 and as high as 90 km above the surface.

   Infrared spectra from NASA's Galileo mission show distinct absorptions
   on Europa, a moon of Jupiter, that have been attributed to one or more
   sulfuric acid hydrates. 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.

History of sulfuric acid

   The discovery of sulfuric acid is credited to the 8th century alchemist
   Jabir ibn Hayyan. It was studied later by the 9th century physician and
   alchemist Ibn Zakariya al-Razi (Rhases), who obtained the substance by
   dry distillation of minerals including iron(II) sulfate heptahydrate,
   FeSO[4] • 7H[2]O, and copper(II) sulfate pentahydrate, CuSO[4] •
   5H[2]O. When heated, these compounds decompose to iron(II) oxide and
   copper(II) oxide, respectively, giving off water and sulfur trioxide,
   which combine to produce a dilute solution of sulfuric acid. This
   method was popularized in Europe through translations of Arabic and
   Persian treatises and books by European alchemists, such as the
   13th-century German Albertus Magnus.

   Sulfuric acid was known to medieval European alchemists as oil of
   vitriol, spirit of vitriol, or simply vitriol, among other names. The
   word vitriol derives from the Latin vitreus, 'glass', for the glassy
   appearance of the sulfate salts, which also carried the name vitriol.
   Salts called by this name included copper(II) sulfate (blue vitriol, or
   rarely Roman vitriol), zinc sulfate (white vitriol), iron(II) sulfate
   (green vitriol), iron(III) sulfate (vitriol of Mars), and cobalt(II)
   sulfate (red vitriol).

   Vitriol was widely considered the most important alchemical substance,
   intended to be used as a philosopher's stone. Highly purified vitriol
   was used as a medium to react substances in. This was largely because
   the acid does not react with gold, often the final aim of alchemical
   processes. The importance of vitriol to alchemy is highlighted in the
   alchemical motto Visita Interiora Terrae Rectificando Invenies Occultum
   Lapidem ('Visit the interior of the earth and rectifying (i.e.
   purifying) you will find the hidden/secret stone'), found in L'Azoth
   des Philosophes by the 15th Century alchemist Basilius Valentinus,
   which is a backronym.

   In the 17th 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 the 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 began producing sulfuric acid this
   way in lead-lined chambers, which were stronger, less expensive, and
   could be made larger than the glass containers which had been used
   previously. This lead chamber process allowed the effective
   industrialization of sulfuric acid production, and with several
   refinements remained the standard method of production for almost two
   centuries.

   John Roebuck's sulfuric acid was only about 35–40% sulfuric acid. Later
   refinements in the lead-chamber process by the French chemist
   Joseph-Louis Gay-Lussac and the British chemist John Glover improved
   this to 78%. However, the manufacture of some dyes and other chemical
   processes require a more concentrated product, and 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, Fe[2](SO[4])[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. The expense of this
   process prevented the large-scale use of concentrated sulfuric acid.

   In 1831, the British vinegar merchant Peregrine Phillips patented a far
   more economical process for producing sulfur trioxide and concentrated
   sulfuric acid, now known as the contact process. Essentially all of the
   world's supply of sulfuric acid is now produced by this method.

Manufacture

   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) + O[2](g) → SO[2](g)

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

          (2) 2 SO[2] + O[2](g) → 2 SO[3](g)     (in presence of V[2]O[5])

   Finally the sulfur trioxide is treated with water (usually as 97-98%
   H[2]SO[4] containing 2-3% water) to produce 98-99% sulfuric acid.

          (3) SO[3](g) + H[2]O( l) → H[2]SO[4](l)

   Note that directly dissolving SO[3] in water is impractical due to the
   highly exothermic nature of the reaction. Mists are formed instead of a
   liquid. Alternatively, the SO[3] is absorbed into H[2]SO[4] to produce
   oleum (H[2]S[2]O[7]), which is then diluted to form sulfuric acid.

          (3) H[2]SO[4]( l) + SO[3] → H[2]S[2]O[7](l)

   Oleum is reacted with water to form concentrated H[2]SO[4].

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

   In 1993, American production of sulfuric acid amounted to 36.4 million
   tonnes. World production in 2001 was 165 million tonnes.

Uses

   Sulfuric acid is a very important commodity chemical, and indeed 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:

          Ca[5]F(PO[4])[3] + 5 H[2]SO[4] + 10 H[2]O → 5 CaSO[4]·2 H[2]O +
          HF + 3 H[3]PO[4]

   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 (SO[2]) and sulfur
   trioxide (SO[3]) 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
   crystalised 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:

          Al[2]O[3] + 3 H[2]SO[4] → Al[2](SO[4])[3] + 3 H[2]O

   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
   H[2]SO[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 important in the manufacture of dyestuffs.

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

Sulfur-iodine cycle

   The sulfur-iodine cycle is a series of thermochemical processes used to
   obtain hydrogen. It consists of three chemical reactions whose net
   reactant is water and whose net products are hydrogen and oxygen.

          2 H[2]SO[4] → 2 SO[2] + 2 H[2]O + O[2] (830°C)
          I[2] + SO[2] + 2 H[2]O → 2 HI + H[2]SO[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. With an efficiency of around 50% it is
   more attractive than electrolysis, and it does not require hydrocarbons
   like current methods of steam reforming. Additionally, the
   sulfur-iodine cycle has a much lower maximum operating temperature
   compared to traditional electrolysis.

   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 large-scale.

Safety

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. The standard first aid treatment for
   acid spills on the skin is, as for other corrosive agents, irrigation
   with large quantities of water: in the case of sulfuric acid it is
   important that the acid should be removed before washing, as a further
   heat burn could result from the exothermic dilution of the acid.
   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.

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 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. In the United States, the permissible exposure limit (PEL) for
   sulfuric acid is fixed at 1 mg/m^3: 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.

In popular culture

In fiction

   The use of sulfuric acid as a weapon in crimes of assault, known as
   "vitriol throwing", has at times been sufficiently common (if
   sensational) to make its way into novels and short stories. Examples
   include The Adventure of the Illustrious Client, by Arthur Conan Doyle,
   and The Love of Long Ago, by Guy de Maupassant. An episode of Saturday
   Night Live hosted by Mel Gibson included a parody Western sketch about
   "Sheriff Jeff Acid," who carries a flask of acid instead of a six
   shooter. The DC Comics villain Two Face was disfigured as a result of a
   vitriol throw. In the novel Veronika Decides to Die by. Paulo Coelho ,
   talks of a girl who has attempted to commit suicide and ends up with
   Vitriol poisoning. Also the docter/therapest in this novel writes a
   thesus on Curing Vitriol poisoning

In comic rhyme

   Sulfuric acid is one of the few compounds whose chemical formula is
   well known by the general public because of many comic rhymes, such as
   this one popular in the UK:

          Johnny was a chemist's son, but Johnny is no more.
          What Johnny thought was H[2]O was H[2]SO[4].

   In the U.S., a common variant is:

          Little Johnny took a drink, but he shall drink no more.
          For what he thought was H[2]O was H[2]SO[4].

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