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The strength of a weak acid is quantified by its acid dissociation constant, 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.
In chemistry, an acid dissociation constant (also known as acidity constant, or acid-ionization constant; denoted ) is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction
The higher the proton affinity, the stronger the base and the weaker the conjugate acid in the gas phase.The (reportedly) strongest known base is the ortho-diethynylbenzene dianion (E pa = 1843 kJ/mol), [3] followed by the methanide anion (E pa = 1743 kJ/mol) and the hydride ion (E pa = 1675 kJ/mol), [4] making methane the weakest proton acid [5] in the gas phase, followed by dihydrogen.
A variety of amines and nitrogen heterocycles are useful bases of moderate strength (pK a of conjugate acid around 10-13) . N,N-Diisopropylethylamine (DIPEA, also called Hünig's Base [1]), pK a = 10.75
In chemistry, biochemistry, and pharmacology, a dissociation constant (K D) is a specific type of equilibrium constant that measures the propensity of a larger object to separate (dissociate) reversibly into smaller components, as when a complex falls apart into its component molecules, or when a salt splits up into its component ions.
The Hammett acidity function (H 0) is a measure of acidity that is used for very concentrated solutions of strong acids, including superacids.It was proposed by the physical organic chemist Louis Plack Hammett [1] [2] and is the best-known acidity function used to extend the measure of Brønsted–Lowry acidity beyond the dilute aqueous solutions for which the pH scale is useful.
Although hydrofluoric acid is regarded as a weak acid, it is very corrosive, even attacking glass when hydrated. [20] Dilute solutions are weakly acidic with an acid ionization constant K a = 6.6 × 10 −4 (or pK a = 3.18), [10] in contrast to corresponding solutions of the other hydrogen halides, which are strong acids (pK a < 0).
The most common Lewis bases are anions. The strength of Lewis basicity correlates with the pK a of the parent acid: acids with high pK a 's give good Lewis bases. As usual, a weaker acid has a stronger conjugate base. Examples of Lewis bases based on the general definition of electron pair donor include: