The term "chromic acid" is usually used for a mixture made by adding concentrated sulfuric acid to a dichromate, which may contain a variety of compounds, including solid chromium trioxide. This kind of chromic acid may be used as a cleaning mixture for glass. Chromic acid may also refer to the molecular species, H2CrO4 of which the trioxide is the anhydride. Chromic acid features chromium in an oxidation state of +6 (and a valence of VI or 6). It is a strong and corrosive oxidizing agent and a moderate carcinogen.
Molecular chromic acid, H2CrO4, in principle, resembles sulfuric acid, H2SO4. It would ionize accordingly:
H2CrO4 ⇌ [HCrO4]− + H+
The pKa for the equilibrium is not well characterized. Reported values vary between about −0.8 to 1.6. The structure of the mono anion has been determined by X-ray crystallography. In this tetrahedral oxyanion, three Cr-O bond lengths are 156 pm and the Cr-OH bond is 201 pm
[HCrO4]− condenses to form dichromate:
2 [HCrO4]− ⇌ [Cr2O7]2− + H2O, logKD = 2.05.
Furthermore, the dichromate can be protonated:
[HCr2O7]− ⇌ [Cr2O7]2− + H+, pKa = 1.8
Loss of the second proton occurs in the pH range 4–8, making the ion [HCrO4]− a weak acid.
Molecular chromic acid could in principle be made by adding chromium trioxide to water (cf.manufacture of sulfuric acid).
CrO3 + H2O ⇌ H2CrO4
In practice, the reverse reaction occurs: molecular chromic acid dehydrates.
Some insights can be gleaned from observations on the reaction of
dichromate solutions with sulfuric acid. The first colour change from
orange to red signals the conversion of dichromate to chromic acid.
Under these conditions deep red crystals of chromium trioxide
precipitate from the mixture, without further colour change.
Structure of tetrachromic acid H2Cr4O13·2H2O,
one component of concentrated "chromic acid". The H-atom positions are
calculated, not observed. Color code: red = O, white = H, blue = Cr.
Higher chromic acids with the formula H2CrnO(3n+1) are probable components of concentrated solutions of chromic acid.
Uses
Chromic
acid is an intermediate in chromium plating, and is also used in
ceramic glazes, and colored glass. Because a solution of chromic acid in
sulfuric acid (also known as a sulfochromic mixture or chromosulfuric acid) is a powerful oxidizing agent, it can be used to cleanlaboratory glassware, particularly of otherwise insoluble organic residues. This application has declined due to environmental concerns. Furthermore, the acid leaves trace amounts of paramagnetic chromic ions (Cr3+) that can interfere with certain applications, such as NMR spectroscopy. This is especially the case for NMR tubes. Piranha solution can be used for the same task, without leaving metallic residues behind.
Chromic acid was widely used in the musical instrument repair industry, due to its ability to "brighten" raw brass. A chromic acid dip leaves behind a bright yellow patina
on the brass. Due to growing health and environmental concerns, many
have discontinued use of this chemical in their repair shops.
It was used in hair dye in the 1940s, under the name Melereon.
It is used as a bleach in processing black and white photographic reversal film.
Reactions
Chromic acid is capable of oxidizing many kinds of organic compounds and many variations on this reagent have been developed:
Oxidation of secondary alcohol to ketone (cyclooctanone) and nortricyclanone.
Use in qualitative organic analysis
In organic chemistry, dilute solutions of chromic acid can be used to oxidize primary or secondary alcohols to the corresponding aldehydes and ketones. Similarly, it can also be used to oxidize an aldehyde to its corresponding carboxylic acid. Tertiary alcohols and ketones are unaffected. Because the oxidation
is signaled by a color change from orange to brownish green (indicating
chromium being reduced from oxidation state +6 to +3), chromic acid is
commonly used as a lab reagent in high school or undergraduate college
chemistry as a qualitative analytical test for the presence of primary
or secondary alcohols, or aldehydes.
Alternative reagents
In oxidations of alcohols or aldehydes into carboxylic acids,
chromic acid is one of several reagents, including several that are
catalytic. For example, nickel(II) salts catalyze oxidations by bleach (hypochlorite). Aldehydes are relatively easily oxidized to carboxylic acids, and mild oxidizing agents
are sufficient. Silver(I) compounds have been used for this purpose.
Each oxidant offers advantages and disadvantages. Instead of using
chemical oxidants, electrochemical oxidation is often possible.
Safety
Hexavalent chromium compounds (including chromium trioxide, chromic acids, chromates, chlorochromates) are toxic and carcinogenic.
Chromium trioxide and chromic acids are strong oxidizers and may react
violently if mixed with easily oxidizable organic substances.
Upon
mixing of concentrated hydrochloric acid and concentrated nitric acid,
chemical reactions occur. These reactions result in the volatile
products nitrosyl chloride and chlorine gas:
HNO3 + 3 HCl → NOCl + Cl2 + 2 H2O
as evidenced by the fuming nature and characteristic yellow color of
aqua regia. As the volatile products escape from solution, aqua regia
loses its potency. Nitrosyl chloride (NOCl) can further decompose into nitric oxide (NO) and elemental chlorine (Cl2):
2 NOCl → 2 NO + Cl2
This dissociation is equilibrium-limited. Therefore, in addition to
nitrosyl chloride and chlorine, the fumes over aqua regia also contain
nitric oxide (NO). Because nitric oxide readily reacts with atmospheric oxygen, the gases produced also contain nitrogen dioxide, NO2 (red fume):
Aqua regia is also used in etching and in specific analytic procedures. It is also used in some laboratories to clean glassware of organic compounds and metal particles. This method is preferred among most over the more traditional chromic acid bath for cleaning NMR tubes, because no traces of paramagnetic chromium can remain to spoil spectra. While chromic acid baths are discouraged because of the high toxicity of chromium
and the potential for explosions, aqua regia is itself very corrosive
and has been implicated in several explosions due to mishandling.
Because its components react quickly, resulting in its decomposition,
aqua regia quickly loses its effectiveness (yet remains a strong acid),
so its components are usually only mixed immediately before use.
Chemistry
Dissolving gold
Pure gold precipitate produced by the aqua regia chemical refining process
Aqua regia dissolves gold,
although neither constituent acid will do so alone. Nitric acid is a
powerful oxidizer, which will dissolve a very small quantity of gold,
forming gold(III) ions (Au3+). The hydrochloric acid provides a ready supply of chloride ions (Cl−), which react with the gold ions to produce tetrachloroaurate(III) anions ([AuCl4]−),
also in solution. The reaction with hydrochloric acid is an equilibrium
reaction that favors formation of tetrachloroaurate(III) anions. This
results in a removal of gold ions from solution and allows further
oxidation of gold to take place. The gold dissolves to become chloroauric acid. In addition, gold may be dissolved by the chlorine present in aqua regia. Appropriate equations are:
Au + 3 HNO 3 + 4 HCl [AuCl 4]− + 3 NO 2 + H 3O+ + 2 H 2O
or
Au + HNO 3 + 4 HCl [AuCl 4]− + NO + H 3O+ + H 2O.
Solid tetrachloroauric acid
may be isolated by evaporating the excess aqua regia, and decomposing
the residual nitric acid by repeatedly heating the solution with
additional hydrochloric acid. That step reduces nitric acid (see decomposition of aqua regia). If elemental gold is desired, it may be selectively reduced with reducing agents such as sulfur dioxide, hydrazine, oxalic acid, etc. The equation for the reduction of oxidized gold (Au3+) by sulfur dioxide (SO2) is the following:
Similar equations can be written for platinum.
As with gold, the oxidation reaction can be written with either nitric
oxide or nitrogen dioxide as the nitrogen oxide product:
The oxidized platinum ion then reacts with chloride ions resulting in the chloroplatinate ion:
Pt4+(aq) + 6 Cl−(aq) → [PtCl6]2−(aq)
Experimental evidence reveals that the reaction of platinum with aqua
regia is considerably more complex. The initial reactions produce a
mixture of chloroplatinous acid (H2[PtCl4]) and nitrosoplatinic chloride ([NO]2[PtCl4]).
The nitrosoplatinic chloride is a solid product. If full dissolution of
the platinum is desired, repeated extractions of the residual solids
with concentrated hydrochloric acid must be performed:
The chloroplatinous acid can be oxidized to chloroplatinic acid by saturating the solution with molecular chlorine (Cl2) while heating:
H2[PtCl4](aq) + Cl2(g) → H2[PtCl6](aq)
Dissolving platinum solids in aqua regia was the mode of discovery for the densest metals, iridium and osmium,
both of which are found in platinum ores and are not dissolved by aqua
regia, instead collecting as insoluble metallic powder (elemental Ir,
Os) on the base of the vessel.
Dissolution of platinum by aqua regia.
Initial state of the transformation.
Intermediate state of the transformation.
Final state of the transformation (four days later).
Precipitating dissolved platinum
As
a practical matter, when platinum group metals are purified through
dissolution in aqua regia, gold (commonly associated with PGMs) is precipitated by treatment with iron(II) chloride. Platinum in the filtrate, as hexachloroplatinate(IV), is converted to ammonium hexachloroplatinate by the addition of ammonium chloride.
This ammonium salt is extremely insoluble, and it can be filtered off.
Ignition (strong heating) converts it to platinum metal:
Unprecipitated hexachloroplatinate(IV) is reduced with elemental zinc, and a similar method is suitable for small scale recovery of platinum from laboratory residues.
Reaction with tin
Aqua regia reacts with tin to form tin(IV) chloride, containing tin in its highest oxidation state:
Aqua regia first appeared in the De inventione veritatis ("On the Discovery of Truth") by pseudo-Geber (after c. 1300), who produced it by adding sal ammoniac (ammonium chloride) to nitric acid.
The preparation of aqua regia by directly mixing hydrochloric acid with
nitric acid only became possible after the discovery in the late
sixteenth century of the process by which free hydrochloric acid can be
produced.
The fox in Basil Valentine's Third Key represents aqua regia, Musaeum Hermeticum, 1678
The third of Basil Valentine'skeys (c. 1600)
shows a dragon in the foreground and a fox eating a rooster in the
background. The rooster symbolizes gold (from its association with
sunrise and the sun's association with gold), and the fox represents
aqua regia. The repetitive dissolving, heating, and redissolving (the
rooster eating the fox eating the rooster) leads to the buildup of
chlorine gas in the flask. The gold then crystallizes in the form of gold(III) chloride, whose red crystals Basil called "the rose of our masters" and "the red dragon's blood". The reaction was not reported again in the chemical literature until 1895.
When Germany invaded Denmark in World War II, Hungarian chemist George de Hevesy dissolved the gold Nobel Prizes of German physicists Max von Laue (1914) and James Franck
(1925) in aqua regia to prevent the Nazis from confiscating them. The
German government had prohibited Germans from accepting or keeping any
Nobel Prize after jailed peace activist Carl von Ossietzky had received the Nobel Peace Prize in 1935. De Hevesy placed the resulting solution on a shelf in his laboratory at the Niels Bohr Institute.
It was subsequently ignored by the Nazis who thought the jar—one of
perhaps hundreds on the shelving—contained common chemicals. After the
war, de Hevesy returned to find the solution undisturbed and precipitated
the gold out of the acid. The gold was returned to the Royal Swedish
Academy of Sciences and the Nobel Foundation. They re-cast the medals
and again presented them to Laue and Franck.
Pure sulfuric acid does not occur naturally due to its strong affinity to water vapor; it is hygroscopic and readily absorbs water vapor from the air.
Concentrated sulfuric acid is a strong oxidant with powerful
dehydrating properties, making it highly corrosive towards other
materials, from rocks to metals. Phosphorus pentoxide is a notable exception in that it is not dehydrated by sulfuric acid but, to the contrary, dehydrates sulfuric acid to sulfur trioxide.
Upon addition of sulfuric acid to water, a considerable amount of heat
is released; thus, the reverse procedure of adding water to the acid is
generally avoided since the heat released may boil the solution,
spraying droplets of hot acid during the process. Upon contact with body
tissue, sulfuric acid can cause severe acidicchemical burns and secondary thermal burns due to dehydration.
Dilute sulfuric acid is substantially less hazardous without the
oxidative and dehydrating properties; though, it is handled with care
for its acidity.
Although nearly 100% sulfuric acid solutions can be made, the subsequent loss of SO3
at the boiling point brings the concentration to 98.3% acid. The 98.3%
grade, which is more stable in storage, is the usual form of what is
described as "concentrated sulfuric acid". Other concentrations are used
for different purposes. Some common concentrations are:
"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.
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, "10 M" 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
Sulfuric acid contains not only H2SO4 molecules, but is actually an equilibrium of many other chemical species, as it is shown in the table below.
Equilibrium of pure sulfuric acid
Species
mMol/kg
HSO−4
15.0
H3SO+4
11.3
H3O+
8.0
HS2O−7
4.4
H2S2O7
3.6
H2O
0.1
Sulfuric acid is a colorless oily liquid, and has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, and 98% sulfuric acid has a vapor pressure of <1 mmHg at 40 °C.
In the solid state, sulfuric acid is a molecular solid that forms monoclinic crystals
with nearly trigonal lattice parameters. The structure consists of
layers parallel to the (010) plane, in which each molecule is connected
by hydrogen bonds to two others. HydratesH2SO4·nH2O are known for n = 1, 2, 3, 4, 6.5, and 8, although most intermediate hydrates are stable against disproportionation.
In spite of the viscosity of the acid, the effective conductivities of the H3SO+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
Acidity
An
experiment that demonstrates the dehydration properties of concentrated
sulfuric acid. When concentrated sulfuric acid comes into contact with sucrose,
slow carbonification of the sucrose takes place. The reaction is
accompanied by the evolution of gaseous products that contribute to the
formation of the foamy carbon pillar that rises above the beaker.Drops of concentrated sulfuric acid rapidly decompose a piece of cotton towel by dehydration.
The product of this ionization is HSO−4, the bisulfate anion. Bisulfate is a far weaker acid:
HSO−4 + H2O → H3O+ + SO2−4 Ka2 = 0.01 (pKa2 = 2)
The product of this second dissociation is SO2−4, the sulfate anion.
Dehydration
Concentrated sulfuric acid has a powerful dehydrating property, removing water (H2O) from other chemical compounds such as table sugar (sucrose) and other carbohydrates, to produce carbon, steam, and heat. Dehydration of table sugar (sucrose) is a common laboratory demonstration. The sugar darkens as carbon is formed, and a rigid column of black, porous carbon called a carbon snake may emerge.
Similarly, mixing starch
into concentrated sulfuric acid gives elemental carbon and water. The
effect of this can also be seen when concentrated sulfuric acid is
spilled on paper. Paper is composed of cellulose, a polysaccharide related to starch. The cellulose reacts to give a burnt appearance in which the carbon appears much like soot that results from fire.
Although less dramatic, the action of the acid on cotton, even in diluted form, destroys the fabric.
The reaction with copper(II) sulfate can also demonstrate the dehydration property of sulfuric acid. The blue crystals change into white powder as water is removed.
Reactions with salts
Sulfuric acid reacts with most bases to give the corresponding sulfate or bisulfate.
Aluminium sulfate, also known as paper maker's alum, is made by treating bauxite with sulfuric acid:
2 AlO(OH) + 3 H2SO4 → Al2(SO4)3 + 4 H2O
Sulfuric acid can also be used to displace weaker acids from their salts. Reaction with sodium acetate, for example, displaces acetic acid, CH3COOH, and forms sodium bisulfate:
Solid state structure of the [D3SO4]+ ion present in [D3SO4]+[SbF6]−, synthesized by using DF in place of HF.
When allowed to react with superacids, sulfuric acid can act as a base and can be protonated, forming the [H3SO4]+ ion. Salts of [H3SO4]+ have been prepared (e.g. trihydroxyoxosulfonium hexafluoroantimonate(V) [H3SO4]+[SbF6]−) using the following reaction in liquid HF:
The above reaction is thermodynamically favored due to the high bond enthalpy of the Si–F bond in the side product. Protonation using simply fluoroantimonic acid, however, has met with failure, as pure sulfuric acid undergoes self-ionization to give [H3O]+ ions:
2 H2SO4 ⇌ H3O+ + HS2O−7
which prevents the conversion of H2SO4 to [H3SO4]+ by the HF/SbF5 system.
The compounds of sulfur and iodine are recovered and reused, hence the process is called the sulfur–iodine 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.
Sulfuric acid is rarely 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.
When sulfur-containing fuels such as coal or oil are burned, sulfur
dioxide is the main byproduct (besides the chief products carbon oxides
and water).
Sulfuric acid is formed naturally by the oxidation of sulfide minerals, such as pyrite:
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 10–50 km above
Earth's surface, sulfuric acid is formed by the oxidation of volcanic
sulfur dioxide by the hydroxyl radical:
SO2 + HO• → HSO3
HSO3 + O2 → SO3 + HO2
SO3 + H2O → H2SO4
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
The permanent Venusian clouds produce a concentrated acid rain, as the clouds in the atmosphere of Earth produce water rain. Sulfuric acid ice has been detected on Jupiter's moon Europa, where it forms when sulfur ions from Jupiter's magnetosphere implant into the icy surface.
In the first step, sulfur is burned to produce sulfur dioxide.
S(s) + O2 → SO2
The sulfur dioxide is oxidized to sulfur trioxide by oxygen in the presence of a vanadium(V) oxidecatalyst. This reaction is reversible and the formation of the sulfur trioxide is exothermic.
2 SO2 + O2 ⇌ 2 SO3
The sulfur trioxide is absorbed into 97–98% H2SO4 to form oleum (H2S2O7),
also known as fuming sulfuric acid or pyrosulphuric acid. The oleum is
then diluted with water to form concentrated sulfuric acid.
H2SO4 + SO3 → H2S2O7
H2S2O7 + H2O → 2 H2SO4
Directly dissolving SO3 in water, called the "wet sulfuric acid process",
is rarely practiced because the reaction is extremely exothermic,
resulting in a hot aerosol of sulfuric acid that requires condensation
and separation.
The sulfur trioxide is hydrated into sulfuric acid H2SO4:
SO3 + H2O → H2SO4(g) (−101 kJ/mol)
The last step is the condensation of the sulfuric acid to liquid 97–98% H2SO4:
H2SO4(g) → H2SO4(l) (−69 kJ/mol)
Other methods
Burning sulfur together with saltpeter (potassium nitrate, KNO3), in the presence of steam, has been used historically. As saltpeter decomposes, it oxidizes the sulfur to SO3, which combines with water to produce sulfuric acid.
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.
A wide variety of laboratory syntheses are known, and typically begin from sulfur dioxide or an equivalent salt. In the metabisulfite method, hydrochloric acid reacts with metabisulfite to produce sulfur dioxide vapors. The 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 conveniently does not produce
an inseparable mist.
3 SO2 + 2 HNO3 + 2 H2O → 3 H2SO4 + 2 NO
Alternatively, dissolving sulfur dioxide in an aqueous solution of an
oxidizing metal salt such as copper(II) or iron(III) chloride:
2 FeCl3 + 2 H2O + SO2 → 2 FeCl2 + H2SO4 + 2 HCl
2 CuCl2 + 2 H2O + SO2 → 2 CuCl + H2SO4 + 2 HCl
Two less well-known laboratory methods of producing sulfuric acid,
albeit in dilute form and requiring some extra effort in purification,
rely on electrolysis. A solution of copper(II) sulfate can be electrolyzed with a copper cathode and platinum/graphite anode to give spongy copper
at cathode and 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):
2 CuSO4 + 2 H2O → 2 Cu + 2 H2SO4 + O2
More costly, dangerous, and troublesome is the electrobromine method, which employs a mixture of sulfur, water, and hydrobromic acid
as the electrolyte. 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 the 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 and hydrogen (omitting losses
of acid as vapors):
Br2 + Br− ↔ Br−3 (initial tribromide production, eventually reverses as Br− depletes)
2 S + Br2 → S2Br2 (bromine reacts with sulfur to form disulfur dibromide)
S2Br2 + 8 H2O + 5 Br2 → 2 H2SO4 + 12 HBr (oxidation and hydration of disulfur dibromide)
Uses
Sulfuric acid production in 2000
Sulfuric acid is a very important commodity chemical, and a nation's
sulfuric acid production was as recently as 2002 believed to be a good
indicator of its industrial strength. World production in the year 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%. World production in 2022 was estimated at about 260 million tonnes.
As of the late 20th century, most of the produced amount (≈60%)
was 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 acidicizing, aluminium
reduction, paper sizing, and water treatment. About 6% of uses are
related to pigments and include paints, enamels,
printing inks, coated fabrics and paper, while 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 dominant use for sulfuric acid is in the "wet method" for the production of phosphoric acid, used for manufacture of phosphatefertilizers.
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:
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.
Sulfuric acid is also important in the manufacture of dyestuffs solutions.
Industrial cleaning agent
Sulfuric acid is used in steelmaking and other metallurgical industries as a pickling agent for removal of rust and fouling. 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 (SO2) and sulfur trioxide (SO3) which are then used to manufacture "new" sulfuric acid.
Hydrogen peroxide (H2O2) can be added to sulfuric acid to produce piranha solution,
a powerful but potentially hazardous 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 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 often used as a dehydrating or
oxidizing agent in industrial reactions, such as the dehydration of
various sugars to form solid carbon.
Electrolyte
Domestic 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):
Sulfuric acid at high concentrations is frequently the major ingredient in domestic acidic drain cleaners which are used to remove lipids, 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.
The study of vitriols (hydrated sulfates of various metals forming glassy minerals from which sulfuric 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 alchemists like the authors writing under the name of Jabir ibn Hayyan (died c. 806 – c. 816 AD, known in Latin as Geber), Abu Bakr al-Razi (865 – 925 AD, known in Latin as Rhazes), Ibn Sina (980 – 1037 AD, known in Latin as Avicenna), and Muhammad ibn Ibrahim al-Watwat (1234 – 1318 AD) included vitriol in their mineral classification lists.
Jabir ibn Hayyan, Abu Bakr al-Razi, Ibn Sina, et al.
The Jabirian authors and al-Razi experimented extensively with the distillation of various substances, including vitriols. In one recipe recorded in his Kitāb al-Asrār ('Book of Secrets'), al-Razi may have created sulfuric acid without being aware of it:
Take white (Yemeni) alum,
dissolve it and purify it by filtration. Then distil (green?) vitriol
with copper-green (the acetate), and mix (the distillate) with the
filtered solution of the purified alum, afterwards let it solidify (or
crystallise) in the glass beaker. You will get the best qalqadis (white
alum) that may be had.
— Abu Bakr al-Razi, Kitāb al-Asrār
In an anonymous Latin work variously attributed to Aristotle (under the title Liber Aristotilis, 'Book of Aristotle'), to al-Razi (under the title Lumen luminum magnum, 'Great Light of Lights'), or to Ibn Sina, the author speaks of an 'oil' (oleum) obtained through the distillation of iron(II) sulfate (green vitriol), which was likely 'oil of vitriol' or sulfuric acid. The work refers multiple times to Jabir ibn Hayyan's Seventy Books (Liber de septuaginta), one of the few Arabic Jabir works that were translated into Latin. The author of the version attributed to al-Razi also refers to the Liber de septuaginta as his own work, showing that he erroneously believed the Liber de septuaginta to be a work by al-Razi. There are several indications that the anonymous work was an original composition in Latin,
although according to one manuscript it was translated by a certain
Raymond of Marseilles, meaning that it may also have been a translation
from the Arabic.
According to Ahmad Y. al-Hassan, three recipes for sulfuric acid occur in an anonymous Garshuni manuscript containing a compilation taken from several authors and dating from before c. 1100 AD. One of them runs as follows:
The water of vitriol and sulphur which is used to
irrigate the drugs: yellow vitriol three parts, yellow sulphur one part,
grind them and distil them in the manner of rose-water.
A recipe for the preparation of sulfuric acid is mentioned in Risālat Jaʿfar al-Sādiq fī ʿilm al-ṣanʿa, an Arabic treatise falsely attributed to the Shi'i Imam Ja'far al-Sadiq (died 765). Julius Ruska dated this treatise to the 13th century, but according to Ahmad Y. al-Hassan it likely dates from an earlier period:
Then distil green vitriol in a cucurbit and alembic,
using medium fire; take what you obtain from the distillate, and you
will find it clear with a greenish tint.
Vincent of Beauvais, Albertus Magnus, and pseudo-Geber
Sulfuric
acid was called 'oil of vitriol' by medieval European alchemists
because it was prepared by roasting iron(II) sulfate or green vitriol in
an iron retort.
The first allusions to it in works that are European in origin appear
in the thirteenth century AD, as for example in the works of Vincent of Beauvais, in the Compositum de Compositis ascribed to Albertus Magnus, and in pseudo-Geber's Summa perfectionis.
Producing sulfuric acid from sulfur
A method of producing oleum sulphuris per campanam,
or "oil of sulfur by the bell", was known by the 16th century: it
involved burning sulfur under a glass bell in moist weather (or, later,
under a moistened bell). However, it was very inefficient (according to Gesner,
5 pounds (2.3 kg) of sulfur converted into less than 1 ounce (0.03 kg)
of acid), and the resulting product was contaminated by sulfurous acid (or rather, solution of sulfur dioxide) so most alchemists (including, for example, Isaac Newton) didn't consider it equivalent with the "oil of vitriol".
In the 17th century, Johann Rudolf Glauber discovered that adding saltpeter (potassium nitrate, KNO3) significantly improves the output, also replacing moisture with steam. As saltpeter decomposes, it oxidizes the sulfur to SO3, 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.
Lead chamber process
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.
Distillation of pyrite
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, FeS2) was heated in air to yield iron(II) sulfate, FeSO4, 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.
Contact process
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.
In the early to mid 19th century "vitriol" plants existed, among other places, in Prestonpans in Scotland, Shropshire and the Lagan Valley in County Antrim, Northern 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.
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.Nitrile glove exposed to drops of 98% sulfuric acid for 10 minutesSuperficial chemical burn caused by two 98% sulfuric acid splashes (forearm skin)
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 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. A saying used to remember
this is "Do like you oughta, add the acid to the water". 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 dangerous, unless the acid is added slowly enough to allow the
mixture sufficient time to cool. Otherwise, the heat produced may be
sufficient to boil the mixture. 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.
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