In chemistry, biochemistry, and pharmacology, a dissociation constant (KD) 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.
One reason for the popularity of the dissociation constant in biochemistry and pharmacology is that in the frequently encountered case where x = y = 1, KD has a simple physical interpretation: when [A] = KD, then [B] = [AB] or, equivalently,
It also presumes the absence of competing reactions, though the derivation can be extended to explicitly allow for and describe competitive binding.
They separate into free and bound components according to the mass conservation principle: To track the concentration of the complex [AB], one substitutes the concentration of the free molecules ([A] or [B]), of the respective conservation equations, by the definition of the dissociation constant, This yields the concentration of the complex related to the concentration of either one of the free molecules Many biological proteins and enzymes can possess more than one binding site.
A simplified mechanism can be formulated if the affinity of all binding sites can be considered independent of the number of ligands bound to the macromolecule.
It can be then assumed that each of these n subunits are identical, symmetric and that they possess only a single binding site.
is defined as the quotient from the portion of bound ligand to the total amount of the macromolecule: K′n are so-called macroscopic or apparent dissociation constants and can result from multiple individual reactions.
For K′3 there are three different dissociation constants — there are only three possibilities for which pocket is filled last (I, II or III) — and one state (I–II–III).
Even when the microscopic dissociation constant is the same for each individual binding event, the macroscopic outcome (K′1, K′2 and K′3) is not equal.
K′3 describes the reaction from three states (two ligands bound) to one state (three ligands bound); hence, the apparent dissociation constant K′3 is three times bigger than the microscopic dissociation constant KD.
The general relationship between both types of dissociation constants for n binding sites is Hence, the ratio of bound ligand to macromolecules becomes where
Then the first equation is proved by applying the binomial rule The dissociation constant is commonly used to describe the affinity between a ligand
Ligand–protein affinities are influenced by non-covalent intermolecular interactions between the two molecules such as hydrogen bonding, electrostatic interactions, hydrophobic and van der Waals forces.
Affinities can also be affected by high concentrations of other macromolecules, which causes macromolecular crowding.
can be described by a two-state process the corresponding dissociation constant is defined where
represent molar concentrations of the protein, ligand, and protein–ligand complex, respectively.
The dissociation constant has molar units (M) and corresponds to the ligand concentration
Sub-picomolar dissociation constants as a result of non-covalent binding interactions between two molecules are rare.
Biotin and avidin bind with a dissociation constant of roughly 10−15 M = 1 fM = 0.000001 nM.
[8] The dissociation constant for a particular ligand–protein interaction can change with solution conditions (e.g., temperature, pH and salt concentration).
This chemical equilibrium is also the ratio of the on-rate (kforward or ka) and off-rate (kback or kd) constants.
Acid dissociation constants are sometimes expressed by pKa, which is defined by This
In this regard, that is depending on the number of the protons they can give up, we define monoprotic, diprotic and triprotic acids.
In the case of multiple pK values they are designated by indices: pK1, pK2, pK3 and so on.
The dissociation constant of water is denoted Kw: The concentration of water, [H2O], is omitted by convention, which means that the value of Kw differs from the value of Keq that would be computed using that concentration.
This variation must be taken into account when making precise measurements of quantities such as pH.