Although there are many different types of corrosive substances in the world, those that are portrayed in films and on television are invariably the strongest acids. This is due to the fact that acids are clearly way more attractive.
They are vaguely used in fictions as just ‘acid’ and are unnamed. So, if we are to figure out which one of the acids is the strongest in the world, we’ll need to be more scientific in measuring the potency of it.
And for those of you who believe pH is the obvious solution, know that this will only get us so far. Although pH does measure how readily an acid becomes negatively charged when it is in water, it only makes sense below pH 0 and is not applicable in non-aqueous systems, where we need to look for the worst of the worst acids.
Hammett Acidity Function
By evaluating how effectively acids can push an electron onto an object other than a water molecule, the Hammett acidity function assigns grades to acids.
Louis Hammett, its creator, examined a variety of reference bases, each one progressively weaker than the last, to determine whether a superacid would completely protonate it, partially protonate it, or not protonate it at all. We can calculate the pH using the equilibrium constant of the reaction between the partially protonated base and acid.
Here is the list that starts with strong acids to the strongest of them all:
10. Sulfuric acid
Sulfuric acid, also referred to as oil of vitriol, is a mineral acid that is well-known for being highly corrosive and one among the strongest acids. Since concentrated sulfuric acid is a strong oxidant with potent dehydrating properties, it is extremely corrosive to other materials, including rocks and metals.
It has a minus (-) 12 Hammett acidity function.
The wet sulfuric acid process or the conventional contact process (DCDA) are both methods for producing sulfuric acid from sulphur, oxygen, and water (WSA).
The production of sulfuric acid, a crucial commodity chemical, is a reliable indicator of a country’s industrial prowess.
Fertilizers use the majority of the sulfuric acid (60%) that is produced annually. In the chemical industry, 20% is utilized in the production of detergents, synthetic resins, dyes, pharmaceuticals, petroleum catalysts, insecticides, among others. The production of explosives, cellophane, acetate and viscose textiles, lubricants, non-ferrous metals, and batteries are just a few of the numerous uses for the remaining 6%.
Sulfuric acid can result in extremely painful burns, especially in high concentrations. Upon contact, it quickly decomposes living tissues like skin and flesh.
If splashed into the eyes, it can cause irreversible blindness by quickly attacking the cornea. It is fatal if consumed and permanently harms internal organs.
Care must be taken when storing sulfuric acid in nonreactive containers like glass.
9. Chlorosulfuric acid
Being the sulfonic acid of chlorine, it is also known as chlorosulfonic acid. It is a distillable, colourless liquid that is lachrymator and hygroscopic. Commercial samples are typically light brown or straw coloured. It is a tetrahedral molecule.
Its Hammett acidity function is minus (-) 12.78.
Chlorosulfates are the salts and esters of chlorosulfuric acid.
Though HSO3Cl is more conventional, the formula is more accurately written as SO2(OH)Cl. It is a chemical and conceptual intermediary between sulfuryl chloride (SO2Cl2) and sulfuric acid (H2SO4).
In an industrial synthesis, sulphur trioxide is dissolved in sulfuric acid and reacts with hydrogen chloride. It can also be made by chlorinating sulfuric acid. The latter approach is more appropriate for lab-scale operations.
Chlorosulfonic acid is used in the synthesis of persulfuric acid and alkyl sulphate. In Ryan Model 147 reconnaissance drones, it has been used as an anti-contrail agent and to create smoke screens.
ClSO3H reacts violently with water to produce sulfuric acid and hydrogen chloride, which are frequently observed as fuming vapours.
Chlorosulfonic acid may be fatal if inhaled, absorbed through the skin or swallowed. It causes eye and skin burns. It is also, water-reactive and a strong oxidizer.
8. Perchloric Acid
Oerchloric acid (HClO4) is a mineral acid. This colourless compound, which is typically found as an aqueous solution, is a stronger acid than sulfuric acid, nitric acid, and hydrochloric acid.
When hot, it is a potent oxidizer, but at room temperature, aqueous solutions up to about 70% by weight are typically safe, exhibiting only strong acid features and no oxidising properties.
Ammonium perchlorate, a crucial component of rocket fuel, can be prepared with the help of perchloric acid, which is also useful for other perchlorate salts.
Perchloric acid is one among the strongest acids and is dangerously corrosive and easily creates mixtures that could explode.
It has a Hammett acidity function of minus (-) 13.
In the middle of the 1810s, Austrian chemist Friedrich von Stadion produced perchloric acid under the name “oxygenated chloric acid.” French pharmacist Georges-Simon Serullas discovered the solid monohydrate of perchloric acid and gave it the modern name.
Manufacturing:
There are two methods for producing perchloric acid industrially. The old method takes advantage of sodium perchlorate’s high aqueous solubility (209 g/100 mL of water at room temperature). Treatment of such solutions with hydrochloric acid gives perchloric acid, precipitating solid sodium chloride.
Perchloric acid is primarily produced as a byproduct of the production of ammonium perchlorate, which is used in rocket fuel. The increased production of perchloric acid has resulted from the growth of rocketry. Every year, several million kilograms are produced.
Perchloric acid has unique properties in analytical chemistry and is a well-proven material for etching liquid crystal displays, crucial electronics applications, and ore extraction. It plays a significant role in the etching of chrome as well.
Perchloric acid is heavily regulated due to its potent oxidising abilities because it can react violently with metals and flammable materials like wood, plastics, and oils.
The O’Connor Plating Works catastrophe which occurred on February 20, 1947, in Los Angeles, California, resulted in 17 fatalities and 150 injuries.
7. Trifluoromethanesulfonic acid
Triflic acid, also known as TFMS, TFSA, HOTf, or TfOH, is the abbreviation for trifluoromethanesulfonic acid, which has the chemical formula CF3SO3H. One of the most strongest acids ever discovered. It is a colourless, hygroscopic liquid that is slightly viscous and soluble in polar solvents.
Triflic acid has a minus (-) 14.1 Hammett acidity function. It falls under the superacid category.
Industrial production of trifluoromethanesulfonic acid involves electrochemical fluorination (ECF) of methanesulfonic acid. The resulting triflate salt is reprotonated after being hydrolyzed from the CF3SO2F. Trifluoromethylsulfenyl chloride, on the other hand, is oxidised to produce trifluoromethanesulfonic acid.
Triflic acid is primarily used in research as an esterification catalyst.
Triflic acid is helpful in protonations in the lab because its conjugate base is nonnucleophilic. Because it behaves as a strong acid in many solvents (acetonitrile, acetic acid, etc.), where common mineral acids (such as HCl or H2SO4) are only moderately strong, it is also used as an acidic titrant in nonaqueous acid-base titration.
Triflate anhydrides, which are potent acylating agents, are produced when triflic acid reacts with acyl halides.
Triflic acid must be added to polar solvents gradually, just like sulfuric acid must, to avoid thermal runaway.
Triflic acid is one of the strongest acids. Contact with skin causes severe burns with delayed tissue destruction. On inhalation it causes fatal spasms, inflammation and edema.
6. Hydrofluoric acid
A mixture of hydrogen fluoride (HF) and water is hydrofluoric acid. HF solutions are odourless, acidic, and extremely corrosive.
It has a Hammett acidity function of minus (-) 15.1.
Hydrofluoric acid was first prepared in 1771, by Carl Wilhelm Scheele. It is now mainly produced by treatment of the mineral fluorite, CaF2, with concentrated sulfuric acid at approximately 265 °C.
The principal use of hydrofluoric acid is in organofluorine chemistry, including PTFE and the widely used antidepressant drug fluoxetine (Prozac) (Teflon). It produces the elemental fluorine. It is commonly used to etch silicon wafers and glass.
Hydrofluoric acid is a highly corrosive liquid that is also a potent contact poison. Because hydrofluoric acid can penetrate tissue, exposure to the skin or eyes, as well as inhalation or ingestion, can easily result in poisoning. The lack of immediate signs of hydrofluoric acid exposure can give victims false hope, leading them to put off seeking medical attention.
5. Fluorosulfuric acid
Fluorosulfuric acid (HSO3F) is an inorganic compound which is one of the most potent acids. It is a tetrahedral molecule that is related to sulfuric acid, H2SO4, by replacing one of the hydroxyl groups with a fluorine atom. Although commercial samples are frequently yellow, it is a colourless liquid.
Although carborane-based acids are stronger, HSO3F is one of the strongest known simple Brnsted acids. It has a minus (-) 15.1 Hammett acidity function.
These acids are all “superacids,” or acids that are stronger than 100% sulfuric acid.
The reaction of HF and sulphur trioxide produces fluorosulfuric acid. KHF2 or CaF2 can also be treated with oleum at 250 °C. HSO3F can be distilled in a glass apparatus after being freed from HF by sweeping with an inert gas.
HSO3F is useful for regenerating HF/H2SO4 mixtures for etching lead glass.
Although it is unclear whether its commercially important, HSO3F isomerizes alkanes and alkylates hydrocarbons with alkenes. It can also be used as a fluorinating agent in the laboratory.
Fluorosulfuric acid is considered extremely toxic and corrosive. It hydrolyzes to produce HF. The addition of water to HSO3F can be violent, similar to but much more violent than the addition of water to sulfuric acid.
4. Carborane superacids
Carborane acids H(CXB11Y5Z6) (X, Y, Z = H, Alk, F, Cl, Br, CF3) are a class of superacids with Hammett acidity function values (H0 -18) and computed pKa values well below -20, making them among the strongest known Brnsted acids.
The highly chlorinated derivative H is the best-studied example (CHB11Cl11). The acidity of H(CHB11Cl11) was discovered to be far greater than that of triflic acid, CF3SO3H, and bistriflimide, (CF 3SO2)2NH, which were previously thought to be the strongest isolable acids.
Their high acidities are result of extensive delocalization of their conjugate bases, carboranate anions (CXB11Y5Z6), which are typically stabilised by electronegative groups such as Cl, F, and CF3. They are the only superacids known to protonate C60 fullerene without decomposing it due to their lack of oxidising properties and the exceptionally low nucleophilicity and high stability of their conjugate bases.
Boron-based carborane acids have numerous proposed applications. They have been proposed as catalysts for hydrocarbon cracking and isomerization of n-alkanes to form branched isoalkanes, for example (“isooctane”, for example).
Carborane acids deliver very clean acidity without ferocity. They are harmful in contact with skin and if inhaled.
3. Magic acid
Magic acid (FSO3HSbF5) is a superacid composed of a 1:1 molar mixture of fluorosulfuric acid (HSO3F) and antimony pentafluoride (SbF5).
The George Olah lab at Case Western Reserve University developed this conjugate Brnsted-Lewis superacid system in the 1960s, and it has been used to stabilise carbocations and hypercoordinated carbonium ions in liquid media.
Magic acid has a minus (-) 19.2 Hammett acidity function.
The term “superacid” was coined in 1927 by James Bryant Conant, who discovered that perchloric acid could protonate ketones and aldehydes in nonaqueous solution to form salts.
Magic acid and other superacids have been shown to protonate even weak bases such as methane, xenon, halogens, and molecular hydrogen when used to catalyse the isomerization of saturated hydrocarbons.
Personal protective equipment, as with all strong acids, especially superacids, should be used. In addition to the required gloves and goggles, a faceshield and full-face respirator are recommended.
Magic acid is highly toxic when ingested or inhaled, causes severe skin and eye burns, and is toxic to aquatic life.
2. Fluoroantimonic acid
Fluoroantimonic acid is a mixture of hydrogen fluoride and antimony pentafluoride that contains a variety of cations and anions (the most basic of which are H2F+ and SbF6).
In terms of protonating ability as measured by the Hammett function (-23<H0<-21), this substance is a superacid that can be over a billion times stronger than 100% pure sulfuric acid. It even protonates some hydrocarbons in order to produce pentacoordinate carbocations (carbonium ions).
This extremely strong acid protonates nearly all organic compounds, frequently resulting in dehydrogenation or dehydration. It is also employed in the production of tetraxenonogold compounds.
Fluoroantimonic acid is extremely corrosive. For example, it cannot be stored directly in glass carboys because it attacks the glass, but it can be stored in PTFE-lined containers (Teflon).
1. Helonium
The helium hydride ion, also known as hydridohelium (1+) ion or helonium, is a positively charged cation with the chemical formula HeH+. It is made up of a helium atom bonded to a hydrogen atom that has had one electron removed. Helonium is also known as protonated helium. It is the lightest heteronuclear ion and is thought to be the first compound formed after the Big Bang.
In 1925, the ion was created in a laboratory for the first time. It is stable in isolation, but extremely reactive, and cannot be prepared in bulk because it will react with any other molecule that comes into contact with it.
Helonium has a minus (-) 63 Hammett acidity function.
The strongest known acid—stronger than even fluoroantimonic acid—its presence in the interstellar medium had been speculated about since the 1970s, and it was finally discovered in April 2019 using the airborne SOFIA telescope.
Because HeH+ cannot be stored in any usable form, the chemistry of it must be studied in situ.
Organic reactions, for example, can be studied by synthesising a tritium derivative of the desired organic compound. The decay of tritium to 3He+, followed by its extraction of a hydrogen atom, produces 3HeH+, which is then surrounded by organic material and reacts.