Acids and bases
Acceptor number Acid Acid–base reaction Acid–base homeostasis Acid strength Acidity function Amphoterism Base Buffer solutions Dissociation constant Donor number Equilibrium chemistry Extraction Hammett acidity function pH Proton affinity Self-ionization of water Titration Lewis acid catalysis Frustrated Lewis pair Chiral Lewis acid
Acid types
Brønsted–Lowry Lewis Mineral Organic Oxide Strong Superacids Weak Solid
Base types
Brønsted–Lowry Lewis Organic Oxide Strong Superbases Non-nucleophilic Weak
v t e

Acid strength is the tendency of an acid, symbolised by the chemical formula ${\displaystyle {\ce {HA}}}$, to dissociate into a proton, ${\displaystyle {\ce {H+}}}$, and an anion, ${\displaystyle {\ce {A-}}}$. The dissociation of a strong acid in solution is effectively complete, except in its most concentrated solutions.

${\displaystyle {\ce {HA -> H+ + A-}}}$

Examples of strong acids are hydrochloric acid ${\displaystyle {\ce {(HCl)}}}$, perchloric acid ${\displaystyle {\ce {(HClO4)}}}$, nitric acid ${\displaystyle {\ce {(HNO3)}}}$ and sulfuric acid ${\displaystyle {\ce {(H2SO4)}}}$.

A weak acid is only partially dissociated, with both the undissociated acid and its dissociation products being present, in solution, in equilibrium with each other.

${\displaystyle {\ce {HA <=> H+ + A-}}}$

Acetic acid (${\displaystyle {\ce {CH3COOH}}}$) is an example of a weak acid. The strength of a weak acid is quantified by its acid dissociation constant, ${\displaystyle K_{{\ce {a}}}}$ value.

The strength of a weak organic acid may depend on substituent effects. The strength of an inorganic acid is dependent on the oxidation state for the atom to which the proton may be attached. Acid strength is solvent-dependent. For example, hydrogen chloride is a strong acid in aqueous solution, but is a weak acid when dissolved in glacial acetic acid.

Measures of acid strength

The usual measure of the strength of an acid is its acid dissociation constant (${\displaystyle K_{{\ce {a}}}}$), which can be determined experimentally by titration methods. Stronger acids have a larger ${\displaystyle K_{{\ce {a}}}}$ and a smaller logarithmic constant (${\displaystyle \mathrm {p} K_{{\ce {a}}}=-\log K_{\text{a}}}$) than weaker acids. The stronger an acid is, the more easily it loses a proton, ${\displaystyle {\ce {H+}}}$. Two key factors that contribute to the ease of deprotonation are the polarity of the ${\displaystyle {\ce {H-A}}}$ bond and the size of atom A, which determine the strength of the ${\displaystyle {\ce {H-A}}}$ bond. Acid strengths also depend on the stability of the conjugate base.

While the ${\displaystyle \mathrm {p} K_{{\ce {a}}}}$ value measures the tendency of an acidic solute to transfer a proton to a standard solvent (most commonly water or DMSO), the tendency of an acidic solvent to transfer a proton to a reference solute (most commonly a weak aniline base) is measured by its Hammett acidity function, the ${\displaystyle H_{0}}$ value. Although these two concepts of acid strength often amount to the same general tendency of a substance to donate a proton, the ${\displaystyle \mathrm {p} K_{{\ce {a}}}}$ and ${\displaystyle H_{0}}$ values are measures of distinct properties and may occasionally diverge. For instance, hydrogen fluoride, whether dissolved in water (${\displaystyle \mathrm {p} K_{{\ce {a}}}}$ = 3.2) or DMSO (${\displaystyle \mathrm {p} K_{{\ce {a}}}}$ = 15), has ${\displaystyle \mathrm {p} K_{{\ce {a}}}}$ values indicating that it undergoes incomplete dissociation in these solvents, making it a weak acid. However, as the rigorously dried, neat acidic medium, hydrogen fluoride has an ${\displaystyle H_{0}}$ value of –15,^[1] making it a more strongly protonating medium than 100% sulfuric acid and thus, by definition, a superacid.^[2] (To prevent ambiguity, in the rest of this article, "strong acid" will, unless otherwise stated, refer to an acid that is strong as measured by its ${\displaystyle \mathrm {p} K_{{\ce {a}}}}$ value (${\displaystyle \mathrm {p} K_{{\ce {a}}}}$ < –1.74). This usage is consistent with the common parlance of most practicing chemists.)

When the acidic medium in question is a dilute aqueous solution, the ${\displaystyle H_{0}}$ is approximately equal to the pH value, which is a negative logarithm of the concentration of aqueous ${\displaystyle {\ce {H+}}}$ in solution. The pH of a simple solution of an acid in water is determined by both ${\displaystyle K_{{\ce {a}}}}$ and the acid concentration. For weak acid solutions, it depends on the degree of dissociation, which may be determined by an equilibrium calculation. For concentrated solutions of acids, especially strong acids for which pH < 0, the ${\displaystyle H_{0}}$ value is a better measure of acidity than the pH.

Strong acids

A strong acid is an acid that dissociates according to the reaction

${\displaystyle {\ce {HA + S <=> SH+ + A-}}}$

where S represents a solvent molecule, such as a molecule of water or dimethyl sulfoxide (DMSO), to such an extent that the concentration of the undissociated species ${\displaystyle {\ce {HA}}}$ is too low to be measured. For practical purposes a strong acid can be said to be completely dissociated. An example of a strong acid is hydrochloric acid.

${\displaystyle {\ce {HCl -> H+ + Cl-}}}$(in aqueous solution)

Any acid with a ${\displaystyle \mathrm {p} K_{{\ce {a}}}}$ value which is less than about -2 is classed as a strong acid. This results from the very high buffer capacity of solutions with a pH value of 1 or less and is known as the leveling effect.^[3]

The following are strong acids in aqueous and dimethyl sulfoxide solution. The values of ${\displaystyle \mathrm{p}{K}_{}}$Zdroj:https://en.wikipedia.org?pojem=Weak_acid
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