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The pH-dependence of the activity displayed by enzymes and the pH-dependence of protein stability, for example, are properties that are determined by the pK a values of amino acid side chains. The p K a values of an amino acid side chain in solution is typically inferred from the p K a values of model compounds (compounds that are similar to ...
Conversely, when pH = pK a, the concentration of HA is equal to the concentration of A −. The buffer region extends over the approximate range pK a ± 2. Buffering is weak outside the range pK a ± 1. At pH ≤ pK a − 2 the substance is said to be fully protonated and at pH ≥ pK a + 2 it is fully dissociated (deprotonated).
The Henderson–Hasselbalch equation can be used to model these equilibria. It is important to maintain this pH of 7.4 to ensure enzymes are able to work optimally. [10] Life threatening Acidosis (a low blood pH resulting in nausea, headaches, and even coma, and convulsions) is due to a lack of functioning of enzymes at a low pH. [10]
The pH of a solution of a monoprotic weak acid can be expressed in terms of the extent of dissociation. After rearranging the expression defining the acid dissociation constant, and putting pH = −log 10 [H +], one obtains pH = pK a – log ( [AH]/[A −] ) This is a form of the Henderson-Hasselbalch equation. It can be deduced from this ...
The smaller the difference, the more the overlap. In the case of citric acid, the overlap is extensive and solutions of citric acid are buffered over the whole range of pH 2.5 to 7.5. Calculation of the pH with a polyprotic acid requires a speciation calculation to be performed. In the case of citric acid, this entails the solution of the two ...
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 pH range is commonly given as zero to 14, but a pH value can be less than 0 for very concentrated strong acids or greater than 14 for very concentrated strong bases. [2] The pH scale is traceable to a set of standard solutions whose pH is established by international agreement. [3]
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 H 0 {\displaystyle H_{0}} value is a better measure of acidity than the pH.