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Sulfuric acid - Sulfuric acid (sulfuric also spelled sulphuric) is a chemical compound of sulfur with the molecular formula H2SO4. It is a colorless, oily, very viscous and hygroscopic liquid. Sulfuric acid is one of the strongest acids and has a strong corrosive effect. This mineral acid forms two series of salts, the hydrogen sulfates and the sulfates, in which one and two protons, respectively, are replaced by cations compared to the free acid. - Sulfuric acid is one of the most technically important chemicals of all and is one of the most widely produced basic chemical substances. In 1993, around 135 million metric tons of sulfuric acid were produced worldwide; in 2020, the figure was 260 million metric tons. It is mainly used in fertilizer production and for the production of other mineral acids, such as hydrochloric or phosphoric acid. Aqueous solutions of various concentrations are mostly used. - The anhydride of sulfuric acid is sulfur trioxide (SO3). The solution of sulfur trioxide in the sulfuric acid beyond the stoichiometric ratio is called fuming sulfuric acid or oleum, because the contained sulfur trioxide easily escapes from the solution and forms mist ("fumes") of dilute sulfuric acid with atmospheric moisture. Related acids are sulfurous acid (H2SO3), which is derived from sulfur dioxide, and thiosulfuric acid (H2S2O3), in which one oxygen atom is replaced by sulfur. – History - Sulfuric acid has been known under the obsolete name vitriol oil for a long time. The first references are found in the texts of the historically disputed alchemist Dschābir ibn Hayyān from the 8th century. After that, possible manufacturing processes are also mentioned in the alchemical writings of Albertus Magnus (1200-1280) and Basilius Valentinus (around 1600). These processes describe how vitriol oil can be obtained from naturally occurring sulfates - such as chalcanthite or alum. The name vitriol oil is derived from the obsolete name vitriol for these minerals. The first source of larger quantities of sulfuric acid was iron vitriol. From the 16th century, sulfuric acid was produced in Bohemia, Saxony and the Harz Mountains using the vitriol process. The product was called Nordhausen vitriol after the first production site in Nordhausen. The first scientific investigations with sulfuric acid were carried out by Johann Rudolph Glauber. He allowed the acid to act on common salt and obtained hydrochloric acid and the Glauber salt sodium sulfate named after him. - However, the processes in which sulfates were used were very complex and expensive. In order to obtain larger quantities, a process was developed in the 18th century in which sulfur and saltpeter (potassium nitrate) were burned in glass vessels. Since the glass vessels were very fragile, the reaction was first carried out in lead containers in 1746 by John Roebuck. Founded in 1778, the Laboratory in Winterthur was the first chemical factory in Switzerland, which producing vitriol oil as its main product. After Nicolas Clément-Desormes and Charles-Bernard Desormes discovered in 1793 that the amount of saltpeter could be significantly reduced by using air, the lead chamber process could be used on a large scale. This was particularly important for the Leblanc process for soda production, invented by Nicolas Leblanc in 1789 and first used by him in 1791. The process was improved several times, for example by the development of methods for absorbing nitrous gases by Joseph Louis Gay-Lussac. It was thus possible to achieve continuous production control. - The main disadvantage of this process was that only a maximum acid concentration of 78 % could be achieved, and more concentrated solutions and oleum still had to be produced by the laborious distillation of iron vitriol. Simple production of higher-concentration sulfuric acid was not possible until Rudolph Messel in England developed the contact process in 1870. – Occurrence - Free sulfuric acid that is not dissociated into oxonium and sulfate ions occurs very rarely in nature. In the atmosphere, it is formed from sulfur dioxide, which is produced during the combustion of sulfur-containing substances or during volcanic eruptions. The sulfur dioxide is oxidized by hydroxyl radicals and oxygen to sulfur trioxide. With water, free sulfuric acid is finally formed. Other oxidizing agents that allow sulfur trioxide to form are ozone or hydrogen peroxide. It then reaches the earth in acid rain in the form of dilute acid. - A small amount of free sulfuric acid also occurs in some volcanic springs called solfataras. - In contrast to the free acid, its salts, especially the sulfates, are much more common in nature. Many different sulfate minerals exist. Among the best known and most important are gypsum (CaSO4·2H2O), barite (BaSO4), chalcanthite (CuSO4·5H2O) or Glauber's salt (Na2SO4·10H2O). - Outside the Earth, sulfuric acid is found in the upper atmosphere of Venus. This is formed by photochemical reactions of sulfur dioxide and water. Droplets are formed that contain 80-85% sulfuric acid. In deeper layers, the high temperatures cause the acid to decompose into sulfur dioxide, oxygen and water, which can rise again to form sulfuric acid. - Extraction and production - The basic material for sulfuric acid production is often elemental sulfur, which is produced in large quantities, 72 million metric tons in 2019, during the desulfurization of natural gas and crude oil and is processed using the Claus process or mined using the Frasch process. The sulfur is burned to obtain sulfur dioxide as a feedstock for the actual production process. - S + O2 → SO2S + O 2 ⟶ SO 2 {\displaystyle {\ce {S + O2 -> SO2}}} - Reaction of sulfur with oxygen - Another source that produces large quantities of sulfur dioxide is the smelting of sulfurous ores. - Examples include copper, zinc or lead extraction from the corresponding sulfides. The sulfur dioxide is formed during roasting with atmospheric oxygen. - 2 ZnS + 3 O2 → 2 ZnO + 2 SO2 - 2 ZnS + 3 O 2 ⟶ 2 ZnO + 2 SO 2 {\displaystyle {\ce {2ZnS + 3O2 -> 2ZnO + 2SO2}}} Reaction during roasting of zinc sulfide - In 1999, three million tons of pyrite were still being roasted in Europe for sulfuric acid production. In Asia, however, the proportion of pyrite is higher. - For resource-poor countries that have neither sulfur nor sulfide ores, the production of sulfuric acid from gypsum (calcium sulfate) using the Müller-Kühne process is an option. In this process, sulfur dioxide is extracted from gypsum and coal in a rotary kiln. The energy-intensive process can be made more profitable if cement is produced as a by-product by adding sand and clay. In the GDR, the process was carried out on a large scale. - For further production, sulfur trioxide must be obtained from the sulfur dioxide. The direct reaction of sulfur and oxygen to form sulfur trioxide takes place only to an insufficient extent, since the equilibrium in the reaction of sulfur dioxide to sulfur trioxide is on the side of sulfur trioxide only at low temperatures. At these temperatures, however, the reaction rate is too low. Therefore, with the aid of suitable catalysts, the reaction must be controlled in such a way that a sufficiently fast reaction is ensured at temperatures that are not too high. - 2 SO 2 + O 2 ↽ − − ⇀ 2 SO 3 {\displaystyle {\ce {2SO2 + O2 <=> 2SO3}}} 2 SO2 + O2 ⇋ 2 SO3 - Reaction of sulfur dioxide to sulfur trioxide - In the contact process, which is still used exclusively, vanadium pentoxide is used as an oxygen-transferring catalyst. A molten salt is formed from vanadium(V) oxide and alkali metal sulfates added as co-catalysts. The reactive complex with the composition [(VO)2O(SO4)4]4− which acts as the actual catalyst, is formed in this melt. Oxygen and sulfur dioxide attach to these without changing the oxidation number of the vanadium and react to form sulfur trioxide. - The temperature during the reaction must be between 420 and 620 °C, since at lower temperatures the catalyst becomes inactive due to the formation of vanadium(IV) compounds and it decomposes at higher temperatures. The reaction is carried out in so-called tray contact furnaces, in which the catalyst is arranged in a total of four layers (the "trays") on top of each other and the gas flowing through is cooled to the appropriate temperature between the trays. - In the so-called double contact process, the sulfur trioxide present is washed out with concentrated sulfuric acid before the last tray. This makes it possible to increase the yield to at least 99.8 % (First General Administrative Regulation on the Federal Immission Control Act, Technical Instructions on Air Quality Control 2002). - After formation of the sulfur trioxide, it is converted to sulfuric acid. For this purpose, remaining sulfur dioxide must first be removed with ammonia or sodium thiosulfate. Since the direct reaction of sulfur trioxide with water is too slow, the gas is passed into concentrated sulfuric acid. In this process, disulfuric acid H2S2O7 is quickly formed. When this is diluted with water, it decomposes to two molecules of sulfuric acid. - SO3 + H2SO4 → H2S2O7 - Conversion of sulfur trioxide with sulfuric acid - H2S2O7 + H2O → 2 H2SO4 - Formation of sulfuric acid - This process does not produce pure sulfuric acid (100%), but concentrated acid with 98% acid content. To produce sulfuric acid 100%, the amount of sulfur trioxide that corresponds to the amount of excess water in the concentrated acid must be introduced into the concentrated acid. - In recent decades, sulfuric acid production has increased sharply, especially in China, while in European countries such as Germany, production has declined. Since the beginning of 2000, China has been dependent on additional quantities from Europe. - Of considerable economic importance in the large-scale industrial production of sulfuric acid is that the three individual steps are exothermic (for values, see Contact process). The heat released is used to generate high-pressure steam for power generation and industrial heating purposes. - Contact process - The contact process is a technical process for the production of sulfuric acid using a catalyst (e.g. vanadium pentoxide on silica). It was used on a large scale, but has since been developed into the more profitable and environmentally friendly double contact process. In the past, the lead chamber process and the vitriol process were also used. - While the lead chamber process dates back to the mid-18th century, the contact process was patented by Peregrine Phillips in Bristol in 1831, but the first industrial implementation did not begin until about 50 years later, a first such plant was built in Freiberg in 1875. Initially, platinum was used as a catalyst, but the real breakthrough came after vanadium was introduced as a catalyst (Chemico 1927). While the lead chamber process still had a share of around 75% in Europe and North America in 1910, it was below 75% in 1930 and only around 15% in 1960, with almost no new plants being built. Today, it has been completely displaced by the contact process. - Process description - In the first step of the process, sulfur dioxide is produced by burning sulfur. The air required for combustion must be sufficiently dried before use to avoid plant corrosion and catalyst deactivation due to sulfuric acid or sulfurous acid that would otherwise be produced: - S8 (s) + 8 O2 (g) → 8 SO2 (g) ; ΔH = -2376 kJ/mol - Sulfur combustion is carried out in excess air in a furnace with refractory lining to produce a gas mixture containing about 10 to 11% sulfur dioxide. After combustion, the gas must be cooled to about 410 to 440 °C to set the temperature for the subsequent catalytic oxidation step. - Sulfur dioxide can also be produced by roasting sulfide ores: - 4 FeS2 + 11 O2 → 2 Fe2O3 + 8 SO2 - The resulting sulfur dioxide is reacted with oxygen in an equilibrium reaction with a platinum or vanadium catalyst (on silica gel SiO2) to form sulfur trioxide: - 2 SO2 (g) + O2 (g) ⇌ 2 SO3 (g) ; ΔH = −198 kJ/mol - The sulfur trioxide obtained reacts with water to form sulfuric acid: - SO3 (g) + H2O (g) → H2SO4 (l) ; ΔH = −177 kJ/mol - SO3 has a higher solubility in H2SO4 than in water. This produces disulfuric acid (also called fuming sulfuric acid or oleum): - SO3 (g) + H2SO4 (l) → H2S2O7 (l) - This can then be mixed with water to obtain twice the amount of sulfuric acid used: - H2S2O7 (l) + H2O (l) → 2H2SO4 (l) It is common practice in most sulfuric acid plants to use approximately 97 to 99% sulfuric acid when dissolving SO3, and to adjust the concentration of this sulfuric acid by adding water so that no fuming sulfuric acid is produced. In some sulfuric acid plants, however, oleum is also deliberately produced, which is then not diluted with water, but used for special applications. - It is important during the reaction of the sulfur dioxide with oxygen to form sulfur trioxide that the temperature does not exceed a range of 400-600 °C. – Catalysis - The essential reaction step is the oxidation of sulfur dioxide with atmospheric oxygen to sulfur trioxide with the aid of vanadium pentoxide as catalyst. Vanadium pentoxide is not contained as a solid in the pores of the diatomaceous earth carrier, but is dissolved in an alkali sulfate melt in the active state. The melting temperature of the alkali sulfate therefore indicates the lower operating limit of the catalyst. More recent catalyst developments lower this melting point, and thus the lower limit of use, by means of cesium doping. - In catalysis, the reactive species is a complex with the composition [(VO)2O(SO4)4]4-. Oxygen first attaches to this, followed by sulfur dioxide. In two stages, a total of two molecules of sulfur dioxide react with the oxygen to form sulfur trioxide. - This sulfur trioxide is introduced into sulfuric acid and H2S2O7 is formed, with water this reacts further to form sulfuric acid. - Double contact process - The double contact process for the production of sulfuric acid is a further development of the contact process, but is more profitable and environmentally compatible and is therefore used on an industrial scale today. - In contrast to the single contact process, the sulfur dioxide is passed over a further contact layer after passing through three contact layers and an intermediate absorber. The resulting sulfur trioxide is then dissolved in sulfuric acid in the final absorber. Modern plants thus achieve a conversion of the sulfur dioxide of at least 99.8 %. – Properties - Physical properties -Anhydrous sulfuric acid is a viscous (cross-linked by H-bonds), colorless liquid with a high density (1.84 g/cm³ at 25 °C), which solidifies below 10.4 °C. The melting point is greatly reduced by small amounts of water and is, for example, 3.0 °C for a 98% sulfuric acid. The frequent light brown coloration of technical sulfuric acid is due to organic impurities that are carbonized by dehydration. Above the boiling point of 279.6 °C of anhydrous sulfuric acid, sulfuric acid vapors form containing excess sulfur trioxide, with the water remaining in the boiling sulfuric acid. The anhydrous sulfuric acid thus passes to a sulfuric acid 98.33% with a constant boiling point of 338 °C. At this temperature, the vapor also has an acid content of 98.33% and thus corresponds to an azeotropic water-sulfuric acid mixture. An acid of the same composition and boiling point is obtained when dilute acid is distilled. Therefore, sulfuric acid 100% cannot be obtained by distilling dilute sulfuric acid, but only by dissolving a certain amount of sulfur trioxide in concentrated sulfuric acid. When heated further above 338 °C, sulfuric acid decomposes into water and sulfur trioxide ("fuming sulfuric acid") and is almost completely dissociated at 450 °C. - As a solid, sulfuric acid crystallizes in the monoclinic crystal system in the space group C2/c (space group no. 15). The lattice parameters are a = 814 pm, b = 470 pm, c = 854 pm and β = 111°. The structure is a wavy layered structure in which each dihydrogensulfate tetrahedron is connected to four other tetrahedra by hydrogen bonds. In addition to crystalline pure sulfuric acid, several sulfuric acid hydrates are known. One example is the dihydrate H2SO4·2H2O, which also crystallizes monoclinically with the space group C2/c (No. 15). A total of six different hydrates with one, two, three, four, six, and eight equivalents of water are known in which the acid is completely split into oxonium and sulfate ions. The oxonium ions are associated with a different number of water molecules depending on the hydrate. The melting point of these hydrates decreases as the number of water molecules increases. Thus, the monohydrate melts at 8.59 °C, while the octahydrate melts at -62 °C. Strong hydrogen bonds act between the individual molecules, which cause the high viscosity of 24.6 mPa·s at 25 °C. In comparison, water has a much lower viscosity of 0.89 mPa·s at 25 °C. Similar to pure water, pure sulfuric acid conducts electricity to a small extent. The specific conductivity is 1.044·10-2 S/cm. The reason for this is the low dissociation of the acid due to autoprotolysis. Dilute acid, on the other hand, conducts electricity well due to the oxonium ions it contains. - 2 H2SO4 ⇋ HSO4- + H3SO4+ - Autoprotolysis reaction - Individual sulfuric acid molecules are present in the gas phase. These have a tetrahedral structure with bond angles of 101.3° between the OH groups and 123.3° between the oxygen atoms. The bond lengths of the sulfur-oxygen bonds are different with 157.4 pm (to OH groups) and 142.2 pm (to oxygen atoms), respectively. The molecular structure in the solid state corresponds to that in the gas phase. - The bonds in the sulfuric acid molecule can be described by various mesomeric boundary structures. For example, the structure in which double bonds are assumed between sulfur and oxygen or in which there are only single bonds and simultaneous charge separation. Theoretical calculations have shown that the 3d orbitals contribute very little to the bonding. Therefore, the real bonding situation in the sulfuric acid molecule is most accurately described by the structure in which only single bonds are drawn. The shortened S-O bond can be explained by additional electrostatic interactions between the charged atoms. - Chemical properties - As a very strong acid, sulfuric acid readily releases protons. With a pKa value of -3.0 (however, this only applies to dilute solutions) or, more precisely, an H0 value of -11.9, sulfuric acid is one of the strong acids in the first protolysis stage. - H2SO4 + H2O → HSO4- + H3O+ - H 2 SO 4 + H 2 O ⟶ HSO 4 − + H 3 O + {\displaystyle {\ce {H2SO4 + H2O -> HSO4- + H3O+}}} Reaction with water in the first protolysis stage - It is not usually counted among the superacids, but it is chosen as the starting point for the definition of superacid: All acids that are stronger than pure sulfuric acid and thus can protonate it are called superacids. - The second protolysis step from hydrogen sulfate to sulfate has a pKa value of 1.9, so the hydrogen sulfate ion is only a medium-strength acid. - HSO 4 − + H 2 O ⟶ SO 4 2 − + H 3 O + {\displaystyle {\ce {HSO4- + H2O -> SO4^{2}- + H3O+}}}HSO4- + H2O → SO42- + H3O+ - Reaction with water in the second protolysis stage - For this reason, hydrogen sulfate is present for the most part in dilute sulfuric acid (concentration about 1 mol/l). The H2SO4 molecule is almost completely dissociated, while the reaction to form sulfate takes place only to a small extent (about 1.3 % at 1 mol/l). Only at higher dilutions are larger amounts of sulfate formed. - Sulfuric acid has a high affinity for water. When acid and water are mixed, various hydrates of the form H2SO4 - n H2O (n = 1-4, 6, 8) are formed under strong heat generation. The strong affinity of sulfuric acid for water is also manifested in its ability to remove hydroxyl groups and protons from organic substances. This removal leaves carbon behind, and the organic substance turns black and carbonized. This effect occurs mainly with substances that contain many hydroxyl groups. Examples are many carbohydrates such as glucose or polysaccharides. Furthermore, the large water affinity can be used for condensation reactions. In these reactions, water is removed from an organic compound without carbonization. An example of this is the synthesis of 2-pyrone. - Another indication of the strong hygroscopicity is that the acid dehydrates itself to a small extent: - 2 H 2 SO 4 ↽ − − ⇀ H 3 O + + HS 2 O 7 − {\displaystyle {\ce {2H2SO4 <=> H3O+ + HS2O7-}}} 2 H2SO4 ⇋ H3O+ + HS2O7- - Self-dehydration of sulfuric acid - Concentrated sulfuric acid has an oxidizing effect and is capable of dissolving more noble metals such as copper, mercury or silver when heated. The sulfuric acid is reduced to sulfur dioxide in the process. In contrast, even pure, base iron is not attacked by passivation of concentrated sulfuric acid. Cu + 2 H2SO4 → CuSO4 + SO2 + 2 H2OCu + 2 H 2 SO 4 ⟶ CuSO 4 + SO 2 + 2 H 2 O {\displaystyle {\ce {Cu + 2H2SO4 -> CuSO4 + SO2 + 2H2O}}} - Dissolving copper in concentrated sulfuric acid - Dilute sulfuric acid, on the other hand, has only a slight oxidative effect, since the reaction to sulfur dioxide is inhibited by the solvent water. Only those metals are oxidized or dissolved which, as base elements, can be oxidized by the reaction of protons to hydrogen. – Uses - Sulfuric acid is used in very large quantities and in many areas. Its production volume - along with that of chlorine - is considered a benchmark for industrial development and the performance level of a country. - It is called by different names depending on its concentration. Between 10% and 20%, it is called dilute sulfuric acid or dilute acid. - Accumulator acid or battery acid has an acid concentration of 33.5 %. These acids remain liquid even below 0 °C. - Sulfuric acid with a content of up to about 70 % is called chamber acid, and up to 80 % is called glover acid. Concentrated sulfuric acid has a content of at least 98.3 % (azeotrope). Dilute acid is produced in large quantities as a waste product in titanium oxide or dye production. - Most of it is consumed in the production of fertilizers. Sulfuric acid is mainly used to obtain phosphate and ammonium sulfate fertilizers. The latter is produced by reacting semi-concentrated sulfuric acid with ammonia. - H 2 SO 4 + 2 NH 3 ⟶ ( NH 4 ) 2 SO 4 {\displaystyle {\ce {H2SO4 + 2NH3 -> (NH4)2SO4}}} H2SO4 + 2 NH3 → (NH4)2SO4 - Reaction of sulfuric acid with ammonia - In the production of phosphate fertilizers, sulfuric acid is required to break down the rock phosphate. The reaction produces superphosphate Ca(H2PO4)2/CaSO4 - Ca 3 ( PO 4 ) 2 + 2 H 2 SO 4 ⟶ Ca ( H 2 PO 4 ) 2 + 2 CaSO 4 {\displaystyle {\ce {Ca3(PO4)2 + 2H2SO4 -> Ca(H2PO4)2 + 2CaSO4}}} Ca3(PO4)2 + 2 H2SO4 → Ca(H2PO4)2 + 2 CaSO4 - Digestion of apatite to superphosphate by semi-concentrated sulfuric acid - In addition to ammonium sulfate, other sulfates are also produced by reacting corresponding salts with sulfuric acid. One example is aluminum sulfate obtained from aluminum hydroxide, which is used in large quantities in the paper industry and as a flocculant in water purification. - Since many ores are soluble in sulfuric acid, it can be used as a digestion agent. Examples include the wet process for producing zinc from zinc oxide and the sulfate process for obtaining the white pigment titanium dioxide. Sulfuric acid can be used to digest not only oxide ores, but also those containing other anions such as fluoride or phosphate. The reaction produces the corresponding acids. This process is relevant for the production of some technically important acids. Examples are hydrofluoric acid from fluorite, phosphoric acid from apatite and hydrochloric acid from halite. As battery acid, sulfuric acid is an important component of the lead accumulator, as used in automobiles as a starter battery. As in the lead accumulator, dilute sulfuric acid also serves as an electrolyte in electrolytic processes. Its advantages over other electrolytes lie in its high conductivity and simultaneously low tendency to reduce. - In organic chemistry, fuming sulfuric acid can be used to insert the sulfonic acid group (sulfonation). This is mainly used to produce surfactants for the detergent industry and dyes. - Another functional group that can be introduced with the help of sulfuric acid is the nitro group. This is done with the aid of so-called nitrating acid, a mixture of sulfuric and nitric acid. This is mainly used for the production of explosives, such as trinitrotoluene or nitroglycerine. This is another reason why, since February 1, 2021, the EU has included sulfuric acid in mixtures containing more than 15% among the restricted precursors for explosives, with the consequence that its use, possession, transfer and sale by and to persons not acting for professional or commercial purposes is prohibited; the professional or commercial purpose must be verified at the time of sale, and suspicious transactions must be reported. - In chemical laboratories, sulfuric acid is among the most commonly used chemicals. Along with hydrochloric acid and nitric acid, it is a widely used strong acid. Among other things, it is used to adjust the pH value, as a catalyst, for example for esterifications, and for fuming in digestions. The strong hydrophilic effect of sulfuric acid is used to dry organic substances and gases in desiccators and washing bottles. - Biological significance - The sulfuric acid formed in the air from sulfur dioxide is a component of acid rain, along with the nitric acid formed from nitrogen oxides. Acid rain can cause a drop in pH, especially in weakly buffered soils and water bodies. One effect of a lower pH is a change in the solubility of some metal ions. For example, aluminum, which is harmful to plants, is more soluble in water at lower pH. Similarly, biologically important ions, such as potassium or magnesium, can be washed out more easily. For these reasons, sulfuric acid is considered a possible cause of forest dieback in the 1980s. Technical measures such as flue gas desulfurization at coal-fired power plants and the introduction of low-sulfur fuels mean that so little sulfur dioxide is released in Germany that rainwater contains significantly less sulfuric acid. - Due to its acidity, sulfuric acid has a toxic effect on fish and other aquatic organisms. Thus, in soft water without buffering capacity, the mean lethal concentration (the LC50 value) for fish is 100-330 mg/L, similar to other mineral acids. - In the tailings of ore mines and open-pit lignite mines, sulfuric acid is formed by a combination of abiotic and microbial oxidation of exposed sulfide-bearing minerals. It is washed out by rainwater and collects as acid mine drainage in residual lakes in which hardly any living organisms are found because of the low pH and high heavy metal content. - Safety instructions - Sulfuric acid has a strong irritating and corrosive effect on skin and mucous membranes. It is capable of destroying living tissue (chemical burn). The mechanisms of action of concentrated and diluted sulfuric acid are clearly distinguishable. In the case of dilute sulfuric acid, the increased proton concentration has a corrosive effect, i.e. the effect is similar to that of other dilute acids. Depending on the concentration, the effect on skin contact is mainly local irritation. It is therefore much less dangerous than concentrated sulfuric acid. Due to its strong hydrophilic effect, sulfuric acid has a charring effect and causes severe damage to the skin and eyes even in small quantities. Slowly healing, painful wounds are formed. Sulfuric acid can also be absorbed from the air via vapors; the MAK value is 0.1 mg/m³, the LC50 value is 510 mg/m³ in rats absorbed by inhalation over four hours. - Since a lot of heat is generated when concentrated sulfuric acid reacts with water, it should only be diluted by pouring it into water and not by adding water to the acid. If water is added to sulfuric acid, it can splash and thus burn bystanders. Since 2021, sulfuric acid in a concentration above 15 percent by weight may no longer be supplied to private individuals. – Detection - Concentrated sulfuric acid is detected by reaction with organic substances. For example, if a piece of wood is immersed in concentrated sulfuric acid, it slowly turns black. It is possible to distinguish dilute and concentrated sulfuric acid by different reactions. The different reactivity of the two acids with base metals, such as zinc or iron, is exploited. While hydrogen is already formed at room temperature in the case of dilute acid, the concentrated acid, which contains almost no free oxonium ions, only reacts when heated to form sulfur dioxide and sulfur. - Since sulfuric acid is dissociated in aqueous solution, it cannot be detected directly in it. Instead, the proton concentration and thus the acidic pH can be determined using suitable indicators or a pH meter. The sulfate ion can be determined, for example, by precipitation as poorly soluble barium sulfate.
