Strong acid

A strong acid is an acid that ionizes completely in an aqueous solution by losing one proton, according to the equation

HA(aq) → H⁺(aq) + A⁻(aq)

For sulfuric acid which is diprotic, the "strong acid" designation refers only to dissociation of the first proton

H₂SO₄(aq) → H⁺(aq) + HSO₄⁻(aq)

More precisely, the acid must be stronger in aqueous solution than hydronium ion, so strong acids are acids with a pK_a < −1.74. An example is HCl for which pK_a = -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 extremely corrosive and can dissolve, among other things, glass and all metals except iridium^{[citation needed]}.

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⁺].

Determining acid strength

The strength of an acid, in comparison to other acids, can be determined without the use of pH calculations by observing the following characteristics:

Electronegativity: The higher the electronegativity of a conjugate base in the same period, the more acidic.
Atomic Radius: With increasing atomic radius, acidity also increases. For example, HCl and HI, both strong acids, ionize 100% in water to become their respective ionic constituents. However, HI is stronger than HCl. This is because the atomic radius of an atom of iodine is much larger than that of a chlorine atom. As a result, the negative charge over the I^- anion is dispersed over a larger electron cloud and its attraction for the proton (H⁺) is not as strong as the same attraction in HCl. Therefore, HI is ionized (deprotonated) more readily.
Charge: The more positively charged a species is, the more acidic (neutral molecules can be stripped of protons more easily than anions, and cations are more acidic than comparable molecules).

Some common strong acids (as defined above)

(Strongest to the weakest)

Hydroiodic acid HI (pK_a = −9.3)^[1]
Hydrobromic acid HBr (pK_a = −8.7)^[1]
Perchloric acid HClO₄ (pK_a ≈ −8)^[4]
Hydrochloric acid HCl (pK_a = −6.3)^[1]
Carbonic acid H₂CO₃
Sulfuric acid H₂SO₄ (first dissociation only, pK_a1 ≈ −3)^[4]
p-Toluenesulfonic acid (pK_a = −2.8) Organic soluble strong acid

Almost strong acids

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"

Hydronium ion H₃O⁺ (pK_a = -1.74). For purposes of simplicity, H₃O⁺ is often replaced in a chemical equation with H⁺. However, a bare proton does not exist free in water but instead is bound to one of the lone pairs of electrons on the H₂O molecule.
Nitric acid HNO₃ (pK_a = -1.64)^[4]
Chloric acid HClO₃ (pK_a = -1.0)^[4]
Some chemists include bromic acid (HBrO₃)^{[citation needed]}, perbromic acid (HBrO₄)^{[citation needed]}, iodic acid (HIO₃)^{[citation needed]}, and periodic acid (HIO₄)^{[citation needed]} as strong acids, although these are not universally accepted.

Extremely strong acids (as protonators)

(Strongest to weakest)

Fluoroantimonic acid H[SbF₆]
Magic acid FSO₃HSbF₅
Carborane superacid H(CHB₁₁Cl₁₁)
Fluorosulfuric acid FSO₃H
Triflic acid CF₃SO₃H

References

^ ^a ^b ^c ^d William L. Jolly "Modern Inorganic Chemistry" (McGraw-Hill, 1984), p.177
^ George A. Olah, et. al. Superacid Chemistry, 2nd ed., Wiley, p. 41.
^ That is, the ability of the carborane superacid to protonate a given base (B) is one million times greater than a solution of sulfuric acid, so that the ratio [BH⁺] / [B] is one million times higher. The relative acidities of strong acids can be evaluated using the Hammett acidity function.
^ ^a ^b ^c ^d Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 171. ISBN 978-0-13-039913-7.

Hill, John W., et al. "General Chemistry." 4th ed. New Jersey: Prentice Hall, 2005.

External links

http://www.cm.utexas.edu/academic/courses/Spring2002/CH301/McDevitt/strong.htm
http://jchemed.chem.wisc.edu/Journal/Issues/2000/Jul/abs849.html
Titration of acids - freeware for data analysis and simulation of potentiometric titration curves
Acids and Bases - definitions

[Jolly-1] William L. Jolly "Modern Inorganic Chemistry" (McGraw-Hill, 1984), p.177

[2] George A. Olah, et. al. Superacid Chemistry, 2nd ed., Wiley, p. 41.

[3] That is, the ability of the carborane superacid to protonate a given base (B) is one million times greater than a solution of sulfuric acid, so that the ratio [BH⁺] / [B] is one million times higher. The relative acidities of strong acids can be evaluated using the Hammett acidity function.

[Housecroft-4] Housecroft, C. E.; Sharpe, A. G. (2004). Inorganic Chemistry (2nd ed.). Prentice Hall. p. 171. ISBN 978-0-13-039913-7.

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