Sulfuric acid

sulfuric acid is an extremely corrosive chemical compound with the formula H₂SO₄. It is the chemical compound that is produced the most in the world, which is why it is used as one of the many measures of the industrial capacity of countries. A large part is used to obtain fertilizers. It is also used for the synthesis of other acids and sulfates and in the petrochemical industry.

It is generally obtained from sulfur dioxide, by oxidation with nitrogen oxides in aqueous solution. Processes are normally carried out afterwards to achieve a higher concentration of the acid. Formerly it was called oil or spirit of vitriol, name by which the sulfate salts from which it was produced were known.

The molecule has a pyramidal structure, with the sulfur atom in the center and the four oxygen atoms at the vertices. The two hydrogen atoms are attached to the oxygen atoms not double bonded to sulfur. Depending on the solution, these hydrogens can dissociate. In water it behaves as a strong acid in its first dissociation, giving the hydrogen sulfate anion, and as a weak acid in the second, giving the sulfate anion.

It has a great dehydrating effect on hydrocarbon molecules such as sucrose. This means that it is capable of capturing its molecules in the form of water, leaving the carbon atoms free with the consequent formation of pure carbon.

History

The discovery of sulfuric acid is related to the 7th century alchemist Jabir ibn Hayyan. It was studied later, in the 9th century, by the alchemist Ibn Zakariya al-Razi, who obtained the substance from the dry distillation of minerals including the mixture of iron(II) sulfate (FeSO₄) with water. and copper(II) sulfate (CuSO₄). On heating, these compounds decompose into iron(II) oxide and copper(II) oxide, respectively, giving water and sulfur trioxide, which combined produces a dilute solution of sulfuric acid. This method became popular in Europe through the translation of Arabic and Persian treatises and books by European alchemists of the 13th century like Saint Albert the Great.

Sulfuric acid was known to alchemists in medieval Europe as oil of vitriol, vitriol liquor, or simply vitriol, among other names. The word vitriol is derived from the Latin "vitreus", meaning 'crystal', and refers to the appearance of sulfate salts, which are also called vitriol. Salts so named include copper(II) sulfate (or 'blue vitriol' or 'Roman vitriol'), zinc sulfate (or 'white vitriol'), iron(II) sulfate (or 'green vitriol'), iron(III) sulfate (or 'Mars vitriol'), and cobalt(II) sulfate (or 'red vitriol').

Vitriol was considered the most important chemical, and it was tried to be used as a philosopher's stone. Highly purified, vitriol was used as a medium to react substances in it.

In the 17th century, the German chemist Johann Glauber obtained sulfuric acid by burning sulfur with potassium nitrate (KNO₃), in the presence of steam. As the potassium nitrate decomposed, the sulfur oxidized to SO₃, which combined with water produced sulfuric acid. In 1736, Joshua Ward, a London pharmacist used this method to begin producing sulfuric acid in large quantities.

In 1746 in Birmingham, John Roebuck began producing it this way in lead chambers, which were stronger and more resistant and cheaper than the glass chambers that had been used before. This lead chamber process allowed for the effective industrialization of sulfuric acid production, which with minor improvements maintained this method of production for at least two centuries.

The acid obtained in this way had a concentration of only 35-40%. Subsequent improvements, carried out by the Frenchman Joseph-Louis Gay-Lussac and the British John Glover managed to increase this figure to 78%. However, the manufacture of some dyes and other chemical products that required a higher concentration in their processes was achieved in the 18th century with the dry distillation of minerals with a technique similar to that of the precursor alchemists. Burning pyrite (iron disulfide) with iron sulfate at 480 °C produced sulfuric acid of any concentration, but this process was tremendously expensive and not profitable for large-scale or industrial production.

In 1831, vinegar seller Peregrine Phillips patented a much cheaper process for making sulfur oxide (VI) and concentrated sulfuric acid, now known as the contact process. Currently, most of the supply of sulfuric acid is obtained by this method.

Physical properties

Degrees of sulfuric acid

Although almost 100% sulfuric acid can be produced, the subsequent loss of SO3 at the boiling point raises the concentration to 98.3% acid. The 98% grade is more stable on 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:

“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% for avoid contamination with nitrosylsulfuric acid) and tower acid the acid recovered from the bottom of the Glover tower.<They are now obsolete as commercial concentrations of sulfuric acid, although they can be prepared in the laboratory from concentrated sulfuric acid if necessary. 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 be rise to 80 °C (176 °F) or more.

Sulfuric acid reacts with its anhydride, SO
3, to form H
2S
2O
7, named acid disulfuric acid, fuming sulfuric acid, disulfuric acid or oleum or, less commonly, Nordhausen acid. Oleum concentrations are expressed in terms of % SO
3. (called % oleum) or as % H
2SO
4 or as % H
2SO
4 (the amount obtained if H
2O; common concentrations en are 40% oleum (109% H
2SO
4) and 65% oleum (114.6% H
2SO
4). Pure H
2S
2O
7 is a solid with a 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, and 98% sulfuric acid has a vapor pressure of <1 mmHg at 40 °C.>.

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

Commercial sulfuric acid is sold in several different purity grades. The technical degree H
2SO
4 is impure and often colored, but it is suitable for making fertilizers. Pure grades, like the United States Pharmacopeia (USP), are used to make pharmaceuticals, dyes, and other things. Analytics scores are also available.

Nine hydrates are known, but three of them were confirmed to be tetrahydrates (H₂SO₄-4H₂O), hemihexahydrate (H₂SO₄-6 ¹⁄₂H₂O) and octahydrate (H₂SO₄-8H₂O).

Polarity and conductivity

The anhydrous H
2SO
4 is a very polar liquid with a dielectric constant of about 100. It has a high electrical conductivity, caused by self-dissociation into a protonated ion and a hydrogen sulfate ion through the process known as autoprotolysis.

2 H
2SO
4 H
3SO+
4 + HSO−
4

The equilibrium constant for autoprotolysis is

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

For comparison, the equilibrium constant for water, K_w is 10^-14, a factor of 10¹⁰ (10 billion) minor.

Despite the viscosity of the acid, the effective conductivities of the ions H
3SO+
4 and HSO−
4 are high due to an intramolecular switching mechanism of protons—analogous to the Grotthuss mechanism in water—so sulfuric acid is a good conductor of electricity. It is also an excellent solvent for many reactions.

Acid formation

Sulfuric acid is commercially available in a large number of concentrations and degrees of purity. There are two main processes for the production of sulfuric acid, the lead chamber method and the contact process. The lead chamber process is the older of the two processes and is currently used to produce much of the acid consumed in the manufacture of fertilizers. This method produces a relatively dilute acid (62-78% H₂SO₄). The contact process produces a purer, more concentrated acid, but requires purer raw materials and the use of expensive catalysts. In both processes sulfur dioxide (SO2) is oxidized and dissolved in water. Sulfur(IV) oxide is obtained by incineration of sulfur, roasting pyrites (iron disulfide), roasting other non-ferrous sulfides, or by burning hydrogen sulfide (H₂S) gas.. Historically there was another method prior to these, but today in disuse, the vitriol process.

Obtained in the laboratory

It can be obtained by passing a current of sulfur dioxide gas (SO₂) in a solution of hydrogen peroxide (H₂O₂ >):

H2O2(aq)+SO2(g)→ → H2SO4(aq){displaystyle {rm {H_{2}O_{2}(aq)+SO_{2}(g)to H_{2}SO_{4}(aq)}}}}}}}}}}

This solution is concentrated by evaporating the water.

Lead chamber process

In the lead chamber process, hot gaseous sulfur dioxide (SO₂) enters through the lower part of a reactor called Glover's tower where it is washed with nitrous vitriol (sulfuric acid with oxide nitric acid (NO) and nitrogen dioxide (NO₂) dissolved in it), and mixed with nitrogen oxide (NO) and nitrogen oxide (IV) (NO₂) gaseous. Part of the sulfur(IV) oxide is oxidized to sulfur(VI) oxide (SO₃) and dissolved in the acid bath to form Glover's or tower's acid (approximately 78% H ₂SO₄).

SO2+NO2Δ Δ NO+SO3{displaystyle mathrm {SO_{2}+NO_{2}longrightarrow NO+SO_{3}}} }

SO3+H2OΔ Δ H2SO4{displaystyle mathrm {SO_{3}+H_{2}Olongrightarrow H_{2}SO_{4}} }

From Glover's tower a mixture of gases (including sulfur oxides (IV) and (VI), oxides of nitrogen, nitrogen, oxygen, and steam) is transferred to a lead-lined chamber where it is treated with more water. The chamber can be a large box-shaped space or a truncated cone-shaped enclosure. Sulfuric acid is formed by a complex series of reactions; it condenses on the walls and accumulates on the floor of the chamber. There can be from three to six chambers in series, where the gases pass through each of the chambers in succession. The acid produced in the chambers, generally called chamber acid or fertilizer acid, contains 62% to 68% H₂SO₄.

NO+NO2+H2OΔ Δ 2HNO2{displaystyle mathrm {NO+NO_{2}+H_{2}Olongrightarrow 2HNO_{2}} }

2HNO2+H2SO3Δ Δ H2SO4+H2O+2NO{displaystyle mathrm {2HNO_{2}+H_{2}SO_{3}longrightarrow H_{2}SO_{4}+H_{2}O+2NO} } }

After the gases have passed through the chambers, they are passed to a reactor called the Gay-Lussac tower where they are washed with cooled concentrated acid (from the Glover tower). Nitrogen oxides and unreacted sulfur dioxide dissolve in the acid to form the nitrous vitriol used in Glover's tower. The remaining gases are usually released into the atmosphere.

Contact process

The process is based on the use of a catalyst to convert SO₂ into SO₃, from which sulfuric acid is obtained by hydration.

2SO2+O2Δ Δ 2SO3{displaystyle mathrm {2SO_{2}+O_{2}longrightarrow 2SO_{3}}} }

SO3+H2OΔ Δ H2SO4{displaystyle mathrm {SO_{3}+H_{2}Olongrightarrow H_{2}SO_{4}} }

In this process, a dry gas mixture containing 7 to 10% SO₂, depending on the source of SO₂ production (lower value corresponds to to plants that roast pyrites and the higher one to those that burn sulfur), and from 11 to 14% of O₂, is preheated and once purified to the maximum, passes to a converter of one or more catalytic beds, usually platinum or vanadium pentoxide (V₂O₅), where SO₃ is formed. Two or more converters are usually used.

The conversion yields from SO₂ to SO₃ in a plant in normal operation range between 96 and 97%, since the initial efficiency is 98% decreases over time. This reduction effect is more pronounced in plants where starting pyrites with a high arsenic content are used, which is not completely eliminated and accompanies the gases that are subjected to catalysis, causing catalyst poisoning. Therefore, at times, the performance can drop to values close to 95%.

In the second converter, the temperature varies between 500 and 600 °C. This is selected to obtain an optimal equilibrium constant with maximum conversion at minimum cost. The residence time of the gases in the converter is approximately 2-4 seconds.

The gases from the catalysis are cooled to approximately 100 °C and pass through an oleum tower, to achieve the partial absorption of SO₃. The off-gases pass through a second tower, where the remaining SO₃ is washed with 98% sulfuric acid. Finally, the non-absorbed gases are discharged into the atmosphere through a stack.

There is a marked difference between the manufacture of SO₂ by burning sulfur and by roasting pyrites, especially if they are arsenic. The dust produced in the roasting process can never be completely removed and, together with impurities, mainly arsenic and antimony, have a significant influence on the overall yield of the plant.

The production of sulfuric acid by combustion of elemental sulfur presents a better energy balance since it does not have to adjust to the rigid purification systems necessarily necessary in pyrite roasting plants.

SO3+H2SO4Δ Δ H2S2O7{displaystyle mathrm {SO_{3}+H_{2}SO_{4}longrightarrow H_{2}S_{2}O_{7}}}} }

H2O+H2S2O7Δ Δ 2H2SO4{displaystyle mathrm {H_{2}O+H_{2}S_{2}O_{7}longrightarrow 2H_{2}SO_{4}}} }

Applications

Sulfur in the form of sulfate is an important source of nutrition for plants. About 60% of the total production of sulfuric acid is used in the manufacture of fertilizers. most of it is used in the production of phosphoric acid, which in turn is used to make fertilizing materials such as superphosphates of lime, which facilitate the uptake of phosphate by plants. Smaller amounts are used to produce ammonium nitrosulfate, a simple nitrogenous fertilizer obtained chemically from the reaction of nitric and sulfuric acids with ammonia. Other important applications are found in petroleum refining, pigment production, steel treatment, extraction of non-ferrous metals, manufacture of explosives, detergents, plastics, and fibers.

Substantial amounts of sulfuric acid are also used as a reaction medium in petrochemical and organic chemical processes involving reactions such as nitrations, condensations and dehydrations. In the petrochemical industry it is used for the refining, alkylation and purification of crude distillates.

In the inorganic chemical industry, sulfuric acid is used in the production of titanium(IV) oxide pigments, hydrochloric acid and hydrofluoric acid, which has replaced chlorofluorocarbons in the refrigeration industry.

In metal processing, sulfuric acid is used for the treatment of steel, copper, uranium and vanadium and in the preparation of electrolytic baths for the purification and silvering of non-ferrous metals.

Some processes in the wood and paper industry require sulfuric acid, as well as some textile processes, chemical fibers, and fur and leather treatment.

Regarding the direct uses, probably the most important use is the sulfide that is incorporated through organic sulfonation, particularly in the production of detergents. Car batteries contain dilute solutions of sulfuric acid (though caution is warranted as even dilute is still highly caustic). Lithium-sulfur batteries are the object of interest, because they store more energy than lithium inone batteries.

Precautions

Nitrile glove exposed to drops of sulfuric acid at 98% for 10 minutes

Preparing an acid solution can be dangerous because of the heat generated in the process. It is vital that the concentrated acid is added to the water (and not the other way around) to take advantage of the high heat capacity of the water and the higher boiling point of the acid. Acid can be heated to over 100°C which would cause the droplet to boil rapidly. If water is added to concentrated acid, acid splashes may occur.