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Henderson-Hasselbalch equation : ウィキペディア英語版
Henderson–Hasselbalch equation
In chemistry, the Henderson–Hasselbalch equation describes the derivation of pH as a measure of acidity (using , the negative log of the acid dissociation constant) in biological and chemical systems. The equation is also useful for estimating the pH of a buffer solution and finding the equilibrium pH in acid-base reactions (it is widely used to calculate the isoelectric point of proteins).
The equation is given by:
:\mathrm = \mathrmK_\mathrm+ \log_ \left ( \frac^-">)}">)} \right )
Here, is the molar concentration of the undissociated weak acid, is the molar concentration (molarity, ''M'') of this acid's conjugate base and is where is the acid dissociation constant, that is:
:\mathrmK_\mathrm = - \log_ (K_\mathrm) = - \log_ \left ( \frac\mathrm^+" TITLE="\mathrm_\mathrm^+">)()}">)} \right ) for the non-specific Brønsted acid-base reaction: \mathrm + \mathrm_\mathrm \rightleftharpoons \mathrm^- + \mathrm_\mathrm^+
In these equations, denotes the ionic form of the relevant acid. Bracketed quantities such as () and () denote the molar concentration of the quantity enclosed.
==For bases==
For the standard base equation:
:\mathrm + \mathrm^ \rightleftharpoons \mathrm^
A second form of the equation, known as the Heylman Equation, expressed in terms of K_\mathrm where K_\mathrm is the base dissociation constant:
:\mathrmK_\mathrm = - \log_ (K_\mathrm) = - \log_ \left ( \frac^-" TITLE="\mathrm\mathrm^-">)()}^-">)} \right )
In analogy to the above equations, the following equation is valid:
:\mathrm = \mathrmK_\mathrm+ \log_ \left ( \frac^+">)}">)} \right )
Where BH+ denotes the conjugate acid of the corresponding base B. Using the properties of these terms at 25 degrees Celsius one can synthesise an equation for pH of basic solutions in terms of p''K''a and pH:
:\mathrm = \mathrmK_\mathrm + \log_ \left(\frac">)}" TITLE="\mathrm^">)}\right)

抄文引用元・出典: フリー百科事典『 ウィキペディア(Wikipedia)
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