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

In chemistry, an acid dissociation constant (also known as acidity constant, or acid-ionization constant; denoted ${\displaystyle K_{a}}$) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction

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

known as dissociation in the context of acid–base reactions. The chemical species HA is an acid that dissociates into A⁻, the conjugate base of the acid and a hydrogen ion, H⁺.^[a] The system is said to be in equilibrium when the concentrations of its components will not change over time, because both forward and backward reactions are occurring at the same rate.^[1]

The dissociation constant is defined by^[b]

${\displaystyle K_{\text{a}}=\mathrm {\frac {}{}} ,}$ or

${\displaystyle \mathrm {p} K_{{\ce {a}}}=-\log _{10}K_{\text{a}}=\log _{10}{\frac {{\ce {}}}{}}}$

where quantities in square brackets represent the concentrations of the species at equilibrium.^[c][2] As a simple example for a weak acid with K_a = 10⁻⁵, log K_a is the exponent which is -5, so that pK_a = 5. And for acetic acid with K_a = 1.8 x 10⁻⁵, pK_a is close to 5. A higher K_a corresponds to a stronger acid which is more dissociated at equilibrium. For the more convenient logarithmic scale, a lower pK_a means a stronger acid.

Theoretical backgroundedit

The acid dissociation constant for an acid is a direct consequence of the underlying thermodynamics of the dissociation reaction; the pK_a value is directly proportional to the standard Gibbs free energy change for the reaction. The value of the pK_a changes with temperature and can be understood qualitatively based on Le Châtelier's principle: when the reaction is endothermic, K_a increases and pK_a decreases with increasing temperature; the opposite is true for exothermic reactions.

The value of pK_a also depends on molecular structure of the acid in many ways. For example, Pauling proposed two rules: one for successive pK_a of polyprotic acids (see Polyprotic acids below), and one to estimate the pK_a of oxyacids based on the number of =O and −OH groups (see Factors that affect pK_a values below). Other structural factors that influence the magnitude of the acid dissociation constant include inductive effects, mesomeric effects, and hydrogen bonding. Hammett type equations have frequently been applied to the estimation of pK_a.^[3][4]

The quantitative behaviour of acids and bases in solution can be understood only if their pK_a values are known. In particular, the pH of a solution can be predicted when the analytical concentration and pK_a values of all acids and bases are known; conversely, it is possible to calculate the equilibrium concentration of the acids and bases in solution when the pH is known. These calculations find application in many different areas of chemistry, biology, medicine, and geology. For example, many compounds used for medication are weak acids or bases, and a knowledge of the pK_a values, together with the octanol-water partition coefficient, can be used for estimating the extent to which the compound enters the blood stream. Acid dissociation constants are also essential in aquatic chemistry and chemical oceanography, where the acidity of water plays a fundamental role. In living organisms, acid–base homeostasis and enzyme kinetics are dependent on the pK_a values of the many acids and bases present in the cell and in the body. In chemistry, a knowledge of pK_a values is necessary for the preparation of buffer solutions and is also a prerequisite for a quantitative understanding of the interaction between acids or bases and metal ions to form complexes. Experimentally, pK_a values can be determined by potentiometric (pH) titration, but for values of pK_a less than about 2 or more than about 11, spectrophotometric or NMR measurements may be required due to practical difficulties with pH measurements.

Definitionsedit

According to Arrhenius's original molecular definition, an acid is a substance that dissociates in aqueous solution, releasing the hydrogen ion H⁺ (a proton):^[5]

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

The equilibrium constant for this dissociation reaction is known as a dissociation constant. The liberated proton combines with a water molecule to give a hydronium (or oxonium) ion H₃O⁺ (naked protons do not exist in solution), and so Arrhenius later proposed that the dissociation should be written as an acid–base reaction:

${\displaystyle {\ce {HA + H2O <=> A- + H3O+}}}$

Brønsted and Lowry generalised this further to a proton exchange reaction:^[6][7][8]

${\displaystyle {\text{acid}}+{\text{base }}{\ce {<=>}}{\text{ conjugate base}}+{\text{conjugate acid}}}$

The acid loses a proton, leaving a conjugate base; the proton is transferred to the base, creating a conjugate acid. For aqueous solutions of an acid HA, the base is water; the conjugate base is A⁻ and the conjugate acid is the hydronium ion. The Brønsted–Lowry definition applies to other solvents, such as dimethyl sulfoxide: the solvent S acts as a base, accepting a proton and forming the conjugate acid SH⁺.

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

In solution chemistry, it is common to use H⁺ as an abbreviation for the solvated hydrogen ion, regardless of the solvent. In aqueous solution H⁺ denotes a solvated hydronium ion rather than a proton.^[9][10]

The designation of an acid or base as "conjugate" depends on the context. The conjugate acid BH⁺ of a base B dissociates according to

${\displaystyle {\ce {BH+ + OH- <=> B + H2O}}}$

which is the reverse of the equilibrium

Zdroj:https://en.wikipedia.org?pojem=Acid_dissociation_constant
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