There are very few strong acids. The strength of an acid refers to its ability or tendency to lose a proton. A strong acid is one that completely dissociates in water; in other words, one mole of a strong acid HA dissolves in water yielding one mole of H+ and one mole of the conjugate base, A−, and none of the protonated acid HA. In contrast a weak acid only partially dissociates and at equilibrium both the acid and the conjugate base are in solution. Examples of strong acids are hydrochloric acid (HCl), hydroiodic acid (HI), hydrobromic acid (HBr), perchloric acid (HClO4), nitric acid (HNO3) and sulfuric acid (H2SO4). In water each of these essentially ionizes 100%. The stronger an acid is, the more easily it loses a proton, H+. Two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths are also often discussed in terms of the stability of the conjugate base.
Stronger acids have a larger Ka and a more negative pKa than weaker acids.
Sulfonic acids, which are organic oxyacids, are a class of strong acids. A common example is toluenesulfonic acid (tosylic acid). Unlike sulfuric acid itself, sulfonic acids can be solids. In fact, polystyrene functionalized into polystyrene sulfonate is a solid strongly acidic plastic that is filterable.
Superacids are acids stronger than 100% sulfuric acid. Examples of superacids are fluoroantimonic acid, magic acid and perchloric acid. Superacids can permanently protonate water to give ionic, crystalline hydronium "salts". They can also quantitatively stabilize carbocations.
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A strong acid is an acid that ionizes completely in an aqueous solution by losing one proton, according to the equation
For sulfuric acid which is diprotic, the "strong acid" designation refers only to dissociation of the first proton
More precisely, the acid must be stronger in aqueous solution than hydronium ion, so strong acids are acids with a pKa < −1.74. An example is HCl for which pKa = -6.3.[1] This generally means that in aqueous solution at standard temperature and pressure, the concentration of hydronium ions is equal to the concentration of strong acid introduced to the solution. While strong acids are generally assumed to be the most corrosive, this is not always true. The carborane superacid H(CHB11Cl11), which is one million times stronger than sulfuric acid,[2][3] is entirely non-corrosive, whereas the weak acid hydrofluoric acid (HF) is corrosive and can dissolve, among other things, glass[4] and all metals except iridium.
In all other acid-water reactions, dissociation is not complete, so will be represented as an equilibrium, not a completed reaction. The typical definition of a weak acid is any acid that does not dissociate completely. The difference separating the acid dissociation constants of strong acids from all other acids is so small that this is a reasonable demarcation.
Due to the complete dissociation of strong acids in aqueous solution, the concentration of hydronium ions in the water is equal to the total concentration (ionized and un-ionized) of the acid introduced to solution: [H+] = [A−] = [HA]total and pH = −log[H+].
The strength of an acid, in comparison to other acids, can be determined without the use of pH calculations by observing the following characteristics:
HA(aq)+ H2O(l) -> H3O+(aq) + A-(aq)
In a strong acid, equilibrium lies far to the right, meaning that almost all of the original HA is dissociated at equilibrium. A strong acid yields a weak conjugate base (A-), so a strong acid is also described as an acid whose conjugate base is a much weaker base than water.[5]
This is a list of strong acids with pKa < -1.74, which is stronger than hydronium ion, from strongest to weakest.
These do not meet the strict criterion of being more acidic than H3O+, although in very dilute solution they dissociate almost completely, so sometimes they are included as "strong acids"
(Strongest to weakest)
Polarity refers to the distribution of electrons in a bond, the region of space between two atomic nuclei where a pair of electrons is shared. When two atoms have roughly the same electronegativity (ability to attract electrons) the electrons are shared evenly and spend equal time on either end of the bond. When there is a significant difference in electronegativities of two bonded atoms, the electrons spend more time near the nucleus of the more electronegative element and an electrical dipole, or separation of charges, occurs, such that there is a partial negative charge localized on the electronegative element and a partial positive charge on the electropositive element. Hydrogen is an electropositive element and accumulates a slightly positive charge when it is bonded to an electronegative element such as oxygen or bromine. As the electron density on hydrogen decreases it is more easily abstracted, in other words, it is more acidic. Moving from left to right across a row on the periodic table elements become more electronegative (excluding the noble gases), and the strength of the binary acid formed by the element increases accordingly:
| Formula | Name | pKa[7] |
|---|---|---|
| HF | hydrofluoric acid | 3.17 |
| H2O | water | 15.7 |
| NH3 | ammonia | 38 |
| CH4 | methane | 48 |
The electronegative element need not be directly bonded to the acidic hydrogen to increase its acidity. An electronegative atom can pull electron density out of an acidic bond through the inductive effect. The electron-withdrawing ability diminishes quickly as the electronegative atom moves away from the acidic bond. The effect is illustrated by the following series of halogenated butanoic acids. Chlorine is more electronegative than bromine and therefore has a stronger effect. The hydrogen atom bonded to the oxygen is the acidic hydrogen. Butanoic acid is a carboxylic acid.
| Structure | Name | pKa[8] |
|---|---|---|
| butanoic acid or butyric acid | ≈4.8 | |
| 4-chlorobutanoic acid | 4.5 | |
| 3-chlorobutanoic acid | ≈4.0 | |
| 2-bromobutanoic acid | 2.93 | |
| 2-chlorobutanoic acid | 2.86 |
As the chlorine atom moves further away from the acidic O—H bond, its effect diminishes. When the chlorine atom is just one carbon removed from the carboxylic acid group the acidity of the compound increases significantly, compared to butanoic acid (a.k.a. butyric acid). However, when the chlorine atom is separated by several bonds the effect is much smaller. Bromine is much more electronegative than either carbon or hydrogen, but not as electronegative as chlorine, so the pKa of 2-bromobutanoic acid is slightly greater than the pKa of 2-chlorobutanoic acid.
The number of electronegative atoms adjacent an acidic bond also affects acid strength. Oxoacids have the general formula HOX where X can be any atom and may or may not share bonds to other atoms. Increasing the number of electronegative atoms or groups on atom X decreases the electron density in the acidic bond, making the loss of the proton easier. Perchloric acid is a very strong acid (pKa ≈ -8) and completely dissociates in water. Its chemical formula is HClO4 and it comprises a central chlorine atom with three chlorine-oxygen double bonds (Cl=O) and one chlorine-oxygen single bond (Cl—O). The singly bonded oxygen bears an extremely acidic hydrogen atom which is easily abstracted. In contrast, chloric acid (HClO3) is a weaker acid, though still quite strong (pKa = -1.0), while chlorous acid (HClO2, pKa = +2.0) and hypochlorous acid (HClO, pKa = +7.53) acids are weak acids.[9]
Carboxylic acids are organic acids that contain an acidic hydroxyl group and a carbonyl (C=O bond). Carboxylic acids can be reduced to the corresponding alcohol; the replacement of an electronegative oxygen atom with two electropositive hydrogens yields a product which is essentially non-acidic. The reduction of acetic acid to ethanol using LiAlH4 (lithium aluminium hydride or LAH) and ether is an example of such a reaction.
The pKa for ethanol is 16, compared to 4.76 for acetic acid.[8][10]
Another factor that contributes to the ability of an acid to lose a proton is the strength of the bond between the acidic hydrogen and the atom that bears it. This, in turn, is dependent on the size of the atoms sharing the bond. For an acid HA, as the size of atom A increases, the strength of the bond decreases, meaning that it is more easily broken, and the strength of the acid increases. Bond strength is a measure of how much energy it takes to break a bond. In other words, it takes less energy to break the bond as atom A grows larger, and the proton is more easily removed by a base. This partially explains why hydrofluoric acid is considered a weak acid while the other hydrohalic acids (HCl, HBr, HI) are strong acids. Although fluorine is more electronegative than the other halogens, its atomic radius is also much smaller, so it shares a stronger bond with hydrogen. Moving down a column on the periodic table atoms become less electronegative but also significantly larger, and the size of the atom tends to dominate its acidity when sharing a bond to hydrogen. Hydrogen sulfide, H2S, is a stronger acid than water, even though oxygen is more electronegative than sulfur. Just as with the halogens, this is because sulfur is larger than oxygen and the H—S bond is more easily broken than the H—O bond.