As an acid sulfuric acid reacts with most bases to give the corresponding what

Sulfuric acid

2007 Schools Wikipedia Selection. Related subjects: Chemical compounds

Physical properties

Forms of sulfuric acid

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 H₂SO₄ 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 SO₃(g) are added to sulfuric acid, H₂S₂O₇ 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 % SO₃ (called % oleum) or as % H₂SO₄ (the amount made if H₂O were added); common concentrations are 40% oleum (109% H₂SO₄) and 65% oleum (114.6% H₂SO₄). Pure H₂S₂O₇ is in fact a solid, melting point 36 °C.

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 an excellent solvent for many reactions. In fact, the equilibrium is more complex than shown above. 100% H₂SO₄ contains the following species at equilibrium (figures shown as mmol per kg solvent): HSO₄ − (15.0), H₃SO₄ + (11.3), H₃O + (8.0), HS₂O₇ − (4.4), H₂S₂O₇ (3.6), H₂O (0.1).

Chemical properties

Reaction with water

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:

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(l) → SnSO₄ + 2 H₂O + SO₂

Environmental aspects

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+ : FeS₂ + 7/2 O₂ + H₂O → Fe 2+ + 2 SO₄ 2- + 2 H +

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 H₂O → 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. CO₂ → CO + O SO₂ + O → SO₃ SO₃ + H₂O → H₂SO₄

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, FeSO₄ • 7H₂O, and copper(II) sulfate pentahydrate, CuSO₄ • 5H₂O. 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, KNO₃), in the presence of steam. As the 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.

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, 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. 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) + 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)

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

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: Ca₅F(PO₄)₃ + 5 H₂SO₄ + 10 H₂O → 5 CaSO₄·2 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 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: 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.

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 H₂SO₄ → 2 SO₂ + 2 H₂O + O₂ (830°C) I₂ + SO₂ + 2 H₂O → 2 HI + H₂SO₄ (120°C) 2 HI → I₂ + H₂ (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

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 H₂O was H₂SO₄.

Sulfuric acid and reactions with it

Chemical properties of sulfuric acid

Sul­fu­ric acid is one of the strong­est diba­sic acids, which has the for­mu­la H₂­SO₄.

As for its phys­i­cal prop­er­ties, sul­fu­ric acid looks like a thick trans­par­ent oily liq­uid with no smell. H₂­SO₄ has found wide use in in­dus­try, and de­pend­ing on the con­cen­tra­tion of sul­fu­ric acid, the so­lu­tion has many dif­fer­ent prop­er­ties and spheres of ap­pli­ca­tion. Sul­fu­ric acid is used in the fol­low­ing cas­es:

His­to­ry of the dis­cov­ery of sul­fu­ric acid

Sul­fu­ric acid has been known to hu­mans since an­cient times, and it was main­ly found in na­ture, for ex­am­ple in vol­canic lakes.

In the 9th cen­tu­ry, the al­chemist from Per­sia Muham­mad Ar-Razzi ob­tained a so­lu­tion of sul­fu­ric acid by the method of burn­ing cop­per and iron sul­fate.

But the method of the Per­sian al­chemist was im­proved four cen­turies lat­er by the Eu­ro­pean sci­en­tist Al­bert Mag­nus.

The mod­ern in­dus­tri­al (con­tact) method of ob­tain­ing sul­fu­ric acid in­volves ox­i­diz­ing sul­fur diox­ide, a gas which forms on the com­bus­tion of sul­fur or sul­fur pyrite. Sul­fur tri­ox­ide forms, and in­ter­acts with wa­ter.

Con­tact sul­fu­ric acid has a con­cen­tra­tion from 92 to 94 per­cent:

Phys­i­cal and physi­co-chem­i­cal prop­er­ties of sul­fu­ric acid

H₂­SO₄ mix­es with wa­ter and SO₃ in all pro­por­tions.

In aque­ous so­lu­tions of H₂­SO₄ hy­drates form of the type H₂­SO₄∙nH₂O

The boil­ing point of sul­fu­ric acid de­pends on the lev­el of con­cen­tra­tion of the so­lu­tion and reach­es its max­i­mum at a con­cen­tra­tion of over 98 %.

The best-known com­pound in in­dus­try is the caus­tic com­pound oleum, which is a so­lu­tion of SO₃ in sul­fu­ric acid.

When the con­cen­tra­tion of sul­fur tri­ox­ide in oleum is high­er, the boil­ing point drops.

Chem­i­cal prop­er­ties of sul­fu­ric acid

When heat­ed, con­cen­trat­ed sul­fu­ric acid is a strong ox­i­diz­er, which can ox­i­dize many met­als, with the ex­cep­tion of:

When con­cen­trat­ed sul­fu­ric acid ox­i­dizes met­als, it can re­duce to H₂S, S and SO₂.

Ac­tive met­al:

8 Al + 15H₂­SO₄(conc.)→4Al₂(SO₄)₃ + 12H₂O + 3H₂S

Met­al of medi­um ac­tiv­i­ty:

2Cr + 4 H2­SO4(conc.)→ Cr2(SO4)3 + 4 H2O + S

Met­al of low ac­tiv­i­ty

2Bi + 6H₂­SO₄(conc.)→ Bi₂(SO₄)₃ + 6H₂O + 3SO₂

With cold con­cen­trat­ed sul­fu­ric acid, such met­als as iron and alu­minum do not re­act, as they are cov­ered with an ox­ide film. This process is called pas­si­va­tion.

Re­ac­tion of sul­fu­ric acid and H₂O

When H₂­SO₄ is mixed with wa­ter an exother­mic process takes place, i.e. a large amount of heat is re­leased and the so­lu­tion may even boil. When con­duct­ing chem­i­cal ex­per­i­ments, one must al­ways add sul­fu­ric acid to wa­ter, not the oth­er way around.

Sul­fu­ric acid is a strong de­hy­drat­ing sub­stance, and con­cen­trat­ed sul­fu­ric acid forces wa­ter out of var­i­ous com­pounds. It is of­ten used as a dry­ing agent.

Re­ac­tion of sul­fu­ric acid and sug­ar

The affin­i­ty of sul­fu­ric acid for wa­ter can be demon­strat­ed by a clas­sic ex­per­i­ment, by mix­ing con­cen­trat­ed and sug­ar, which is an or­gan­ic com­pound – a car­bo­hy­drate. In or­der to re­move wa­ter from a sub­stance, sul­fu­ric acid is ca­pa­ble of de­stroy­ing mol­e­cules.

To con­duct the ex­per­i­ment, add a few drops of wa­ter to sug­ar and mix, then care­ful­ly add sul­fu­ric acid. Af­ter a short pe­ri­od, one can ob­serve a vi­o­lent re­ac­tion, with the for­ma­tion of car­bon and the re­lease of gas­es, sul­fur diox­ide and car­bon diox­ide.

Sul­fu­ric acid and sug­ar cube:

Re­mem­ber that work­ing with sul­fu­ric acid is very dan­ger­ous with­out ob­serv­ing safe­ty rules, as sul­fu­ric acid is a caus­tic sub­stance that can in­stant­ly leave se­ri­ous burns on the skin.

Here you’ll find safe chem­i­cal ex­per­i­ments with sug­ar

Re­ac­tion of sul­fu­ric acid and zinc

This re­ac­tion is quite pop­u­lar, and are one of the most wide­spread lab­o­ra­to­ry meth­ods for ob­tain­ing hy­dro­gen: if you add zinc gran­ules to di­lut­ed sul­fu­ric acid, the met­al will dis­solve with the re­lease of gas:

Zn + H₂­SO₄ → Zn­SO₄ + H₂

Di­lut­ed sul­fu­ric acid re­acts with met­als which are to the left of hy­dro­gen in the row of ac­tiv­i­ty, ac­cord­ing to the gen­er­al scheme:

Ме + H₂­SO₄(di­lut­ed) → salt + H₂↑

Re­ac­tion of sul­fu­ric acid and bar­i­um

The qual­i­ta­tive re­ac­tion to sul­fu­ric acid and its salt is the re­ac­tion with bar­i­um ions. This re­ac­tion is wide­ly used in quan­ti­ta­tive anal­y­sis, in par­tic­u­lar in gravime­try.

H₂­SO₄ + Ba­Cl₂ → Ba­SO₄ + 2HCl

Zn­SO₄ + Ba­Cl₂ → Ba­SO₄ + Zn­Cl₂

Warn­ing! Don’t try to re­peat these ex­per­i­ments with­out a pro­fes­sion­al su­per­vi­sion!

Sulfuric acid

Sulfuric acid (or sulphuric acid in British English) is a strong mineral acid with the chemical formula H₂SO₄. It is soluble in water at all concentrations. It was once known as oil of vitriol, a term coined by the eighth-century alchemist Jabir ibn Hayyan (Geber), the chemical’s probable discoverer. [1]

Contents

History of sulfuric acid

The discovery of sulfuric acid is credited to the eighth-century alchemist Jabir ibn Hayyan (Geber). It was studied later by the ninth-century physician and alchemist ibn Zakariya al-Razi (Rhases), who obtained the substance by the dry distillation of minerals, including iron(II) sulfate heptahydrate (FeSO₄ • 7H₂O) and copper(II) sulfate pentahydrate (CuSO₄ • 5H₂O). When heated, these compounds decompose to iron(II) oxide and copper(II) oxide, respectively, giving off water and sulfur trioxide. The combination of water with sulfur trioxide produced a dilute solution of sulfuric acid. This method was popularized in Europe through translations of Arabic and Persian treatises and books by European alchemists, including the thirteenth-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 (meaning «glass»), for the glassy appearance of the sulfate salts, which also carried the name vitriol. Salts that were given this name included copper(II) sulfate (blue vitriol, or occasionally 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, a backronym, [2] 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’). [3]

In the seventeenth century, the German-Dutch chemist Johann Glauber prepared sulfuric acid by burning sulfur together with saltpeter (potassium nitrate, KNO₃), in the presence of steam. As the 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.

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 that 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.

Roebuck’s sulfuric acid was only about 35–40 percent 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 percent. However, the manufacture of some dyes and other chemical processes require a more concentrated product, and throughout the eighteenth 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). When iron(III) sulfate was heated to 480 °C, it 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.

Sulfuric acid at various concentrations

Although nearly 100 percent sulfuric acid can be made, it loses sulfur trioxide (SO₃) gas at the boiling point to produce 98.3 percent acid. The 98 percent grade is 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 noted below.

Given that sulfuric acid is a strong acid, a 0.50 Molar (M) solution of this acid has a pH close to zero.

Different purities are also available. Technical grade H₂SO₄ 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 SO₃(g) are added to sulfuric acid, H₂S₂O₇ is formed. It is called pyrosulfuric acid, fuming sulfuric acid, or oleum. A less common name is Nordhausen acid. Concentrations of oleum are expressed either in terms of percent SO₃ (called percent oleum) or percent H₂SO₄ (the amount made if H₂O were added). Common concentrations are 40 percent oleum (109 percent H₂SO₄) and 65 percent oleum (114.6 percent H₂SO₄). Pure H₂S₂O₇ is a solid, with a melting point of 36 °C.

Physical properties

Anhydrous H₂SO₄ is a very polar liquid, with a dielectric constant of around 100. This property arises from the fact that it can dissociate by protonating itself, a process known as autoprotolysis. [4] This protonation occurs to a high degree, more than ten billion times the level seen in water. It may be represented as follows:

This process allows protons to be highly mobile in H₂SO₄. It also makes sulfuric acid an excellent solvent for many reactions. In fact, the chemical equilibrium is more complex than that shown above. At equilibrium, 100 percent H₂SO₄ contains the following species (figures in parentheses indicate amounts in terms of moles per kilogram solvent):

Chemical properties

Reaction with water

The reaction of sulfuric acid with water (called a hydration reaction) produces a large amount of heat, and it is therefore called an exothermic reaction. If water is added to concentrated sulfuric acid, it can boil and spit dangerously. One reason for this behavior is related to the relative densities of the two liquids. Water is less dense than sulfuric acid and will tend to float above the acid.

To dilute the acid safely, one should always add the acid to the water (in small increments) rather than the water to the acid. [5]

The reaction is best thought of as forming hydronium ions, as follows:

Because the hydration of sulfuric acid is thermodynamically favorable, [6] 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 remove hydrogen and oxygen atoms from other compounds. For example, mixing starch (C₆H₁₂O₆)_n and concentrated sulfuric acid will give elemental carbon and water, which is absorbed by the sulfuric acid (which becomes slightly diluted):

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 sulfates. 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:

Sulfuric acid can be used to displace weaker acids from their salts. For example its reaction with sodium acetate gives acetic acid:

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

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.

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

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

Note that directly dissolving SO₃ in water is impractical, because the reaction is highly exothermic and generates mists instead of a liquid.

An alternative method is to absorb SO₃ into H₂SO₄, to produce oleum (H₂S₂O₇). The oleum is then diluted to form sulfuric acid.

Oleum is reacted with water to form concentrated H₂SO₄.

Sulfuric acid is a very important commodity chemical, and indeed a nation’s sulfuric acid production is a good indicator of its industrial strength. [7] The major use (60 percent of total worldwide) for sulfuric acid is in the «wet method» for the production of phosphoric acid, used for manufacture of phosphate fertilizers and trisodium phosphate for detergents. This method involves the use of phosphate rock, and more than 100 million metric tons is processed annually. This raw material, shown below as fluorapatite (Ca₅F(PO₄)₃) (though the exact composition may vary), is treated with 93 percent sulfuric acid to produce calcium sulfate (CaSO₄), hydrogen fluoride (HF), and phosphoric acid (H₃PO₄). The HF is removed as hydrofluoric acid. The overall process can be represented as follows:

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 crystallized out as a salt (often brown because of iron contamination) and sold into the agro-chemicals industry.

Another important use for sulfuric acid is for the manufacture of aluminium sulfate, also known as papermaker’s alum. This can react with small amounts of soap on paper pulp fibers to give gelatinous aluminium carboxylates, which help to coagulate the pulp fibers into a hard paper surface. It is also used for making aluminum hydroxide, which is used at water treatment plants to filter out impurities, as well as to improve the taste of the water. Aluminum sulfate is made by reacting bauxite with sulfuric acid:

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.

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.

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. It does not require hydrocarbons like current methods of steam reforming.

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.

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 sulfur 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+ :

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

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 Venerian atmosphere, the result is sulfur trioxide, which can combine with water vapor, another trace component of Venus’ atmosphere, to yield sulfuric acid.

In the upper, cooler portions of Venus’s 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.

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. one M, ten percent) 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: 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 six M (35 percent) 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 nonflammable, 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 one 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.

See also

Notes

References

External links

All links retrieved January 5, 2020.

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acid–base reaction

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Acids are substances that contain one or more hydrogen atoms that, in solution, are released as positively charged hydrogen ions. An acid in a water solution tastes sour, changes the colour of blue litmus paper to red, reacts with some metals (e.g., iron) to liberate hydrogen, reacts with bases to form salts, and promotes certain chemical reactions (acid catalysis). Bases are substances that taste bitter and change the colour of red litmus paper to blue. Bases react with acids to form salts and promote certain chemical reactions (base catalysis).

Acids and bases are assigned a value between 0 and 14, the pH value, according to their relative strengths. Pure water, which is neutral, has a pH of 7. A solution with a pH less than 7 is considered acidic, and a solution with a pH greater than 7 is considered basic, or alkaline. Strong acids have a higher concentration of hydrogen ions, and they are assigned values closer to 0. Conversely, strong bases have higher concentrations of hydroxide ions, and they are assigned values closer to 14. Weaker acids and bases are closer to the pH value of 7 than their stronger counterparts.

Different reactions produce different results. Reactions between strong acids and strong bases decompose more completely into hydrogen ions (protons, positively charged ions) and anions (negatively charged ions) in water. For a weak acid and a weak base, neutralization is more appropriately considered to involve direct proton transfer from the acid to the base. If one of the reactants is present in great excess, the reaction can produce a salt (or its solution), which can be acidic, basic, or neutral depending on the strength of the acids and bases reacting with one another.

Acids are chemical compounds that show, in water solution, a sharp taste, a corrosive action on metals, and the ability to turn certain blue vegetable dyes red. Bases are chemical compounds that, in solution, are soapy to the touch and turn red vegetable dyes blue. When mixed, acids and bases neutralize one another and produce salts, substances with a salty taste and none of the characteristic properties of either acids or bases.

The idea that some substances are acids whereas others are bases is almost as old as chemistry, and the terms acid, base, and salt occur very early in the writings of the medieval alchemists. Acids were probably the first of these to be recognized, apparently because of their sour taste. The English word acid, the French acide, the German Säure, and the Russian kislota are all derived from words meaning sour (Latin acidus, German sauer, Old Norse sūur, and Russian kisly). Other properties associated at an early date with acids were their solvent, or corrosive, action; their effect on vegetable dyes; and the effervescence resulting when they were applied to chalk (production of bubbles of carbon dioxide gas). Bases (or alkalies) were characterized mainly by their ability to neutralize acids and form salts, the latter being typified rather loosely as crystalline substances soluble in water and having a saline taste.

In spite of their imprecise nature, these ideas served to correlate a considerable range of qualitative observations, and many of the commonest chemical materials that early chemists encountered could be classified as acids (hydrochloric, sulfuric, nitric, and carbonic acids), bases (soda, potash, lime, ammonia), or salts (common salt, sal ammoniac, saltpetre, alum, borax). The absence of any apparent physical basis for the phenomena concerned made it difficult to make quantitative progress in understanding acid–base behaviour, but the ability of a fixed quantity of acid to neutralize a fixed quantity of base was one of the earliest examples of chemical equivalence: the idea that a certain measure of one substance is in some chemical sense equal to a different amount of a second substance. In addition, it was found quite early that one acid could be displaced from a salt with another acid, and this made it possible to arrange acids in an approximate order of strength. It also soon became clear that many of these displacements could take place in either direction according to experimental conditions. This phenomenon suggested that acid–base reactions are reversible—that is, that the products of the reaction can interact to regenerate the starting material. It also introduced the concept of equilibrium to acid–base chemistry: this concept states that reversible chemical reactions reach a point of balance, or equilibrium, at which the starting materials and the products are each regenerated by one of the two reactions as rapidly as they are consumed by the other.

Acids, Bases and Salts

We have a good understanding of acids and bases in modern chemistry (also called alkalis). Acids and bases are utilized as laboratory reagents, industrial catalysts, culinary additives, and cleaning products, and they pervade our life from the laboratory to the kitchen. However, it took centuries for chemists to completely comprehend these chemicals over the course of history.

What are Acids?

An acid is a molecule that can contribute an H+ ion while also remaining energetically favourable after losing that ion. e.g. Sulfuric acid (H₂SO₄), Acetic Acid (CH₃COOH), Nitric Acid (HNO₃) etc.

Following are some physical properties of acids:

Following are some chemical properties of acids:

Metal + Acid → Salt + Hydrogen

e.g. When hydrochloride acid combines with zinc metal, it produces hydrogen gas and zinc chloride.

Metal carbonate + Acid → Salt + Carbon dioxide + Water

e.g. When hydrochloric acid combines with sodium carbonate, it produces carbon dioxide gas, sodium chloride, and water.

Acid + Metal hydrogen carbonate → Salt + Carbon dioxide + Water

e.g. Sulfuric acid gives sodium sulfate, Carbon dioxide gas and water when it reacts with sodium bicarbonate.

Types of Acids

Acids are classified on different bases like:

Applications of Acids:

What are Bases?

The term “alkali” refers to a base that can be dissolved in water. When these compounds react chemically with acids, they produce salts and hydroxide ions (OH – ) in water. For example Potassium hydroxide (caustic potash or KOH), Calcium hydroxide (Ca(OH)₂), Sodium hydroxide (caustic soda or NaOH) etc.

Following are some physical properties of bases:

Following are some chemical properties of bases:

Alkali + Metal → Salt + Hydrogen

For Example: When sodium hydroxide interacts with aluminium metal, sodium aluminate and hydrogen gas are generated.

Non-metallic oxide + Base → Salt + Water

Alkali + Ammonium salt → Salt + Water + Ammonia

When calcium hydroxide reacts with ammonium chloride, calcium chloride, water, and ammonia are produced.

Types of Bases

Acidity, concentration, and degree of ionization are three variables that can be used to classify bases.

Applications of Bases:

What are Salts?

When an acid and a base react to neutralise one another, they generate sales, which are ionic substances. Salts do not have an electrical charge. Salts come in a variety of forms, the most common of which being sodium chloride. Table salt or common salt are both terms for sodium chloride. Sodium chloride is used to make dishes taste better.

Following are some physical properties of Salts:

Types of Salts

Neutral, Acidic and Basic Salts

Neutral Salts: Salts generated by the reaction of a strong acid with a strong base are neutral in nature. The pH of these salts is 7, which is considered neutral. Potassium chloride, sodium chloride, and sodium sulphate

NaOH + HCl → NaCl + H₂O

Acidic Salts: Acidic salts are the salts formed when a strong acid reacts with a weak base. The pH of acidic salt is less than 7. Examples include ammonium sulphate, ammonium chloride, and various ammonium compounds.

Basic Salts: The salts formed when a weak acid reacts with a strong base are known as basic salts. Examples include sodium carbonate, sodium acetate, and other salts.

What is the cause of the formation of acidic, basic and neutral salts?

Sample Questions

Question 1: When HCl combines with salt, what happens?

Answer:

The acid is dilute hydrochloric acid, and the metal is iron in this example. To make iron (II) chloride and hydrogen, dilute hydrochloric acid is poured to the iron filings. Iron substitutes hydrogen from hydrochloric acid in this process, resulting in iron chloride and hydrogen. This is a basic displacement reaction for gas.

Question 2: What is the most important distinction between an acid and a base?

Answer:

Acids and bases are two types of corrosive chemicals. Acidic materials have a pH value between 0 and 7, while bases have a pH value between 7 and 14. Acids are ionic chemicals that break down in water to create the hydrogen ion (H+).

Question 3: What are the physical properties of bases?

Answer:

Question 4: When the pH of the mouth falls below 5.5, why does tooth decay begin?

Answer:

When the pH of our mouth falls below 5.5, tooth decay begins. This is because below this pH value, the mouth’s medium becomes more acidic, causing tooth enamel to deteriorate more quickly.

Question 5: Write the physical properties of acids.

Answer:

The properties of the acids are:

Question 6: What will happen when sodium hydroxide interacts with aluminium metal?

Answer:

When sodium hydroxide interacts with aluminium metal, sodium aluminate and hydrogen gas are generated

Question 7: What will happen when hydrochloric acid combines with sodium carbonate?

Answer:

When hydrochloric acid combines with sodium carbonate, it produces carbon dioxide gas, sodium chloride, and water.