VPB 112: General Veterinary Biochemistry (1+1)

HENDERSON - HASSELBALCH EQUATION

• A weak acid HA ionizes as follows

 HA → H⁺ + A^-

• The equilibrium constant for this reaction is written as follows

K_a = [H⁺][A^-] / [HA]

The value of te dissociation constant (Ka) indicates the tendency of the acid to lose its proton. Stronger acids have a grater tendency to dissociate and therefore have higher dissoiation constant

• [HA] = concentration of the undissociated acid, [H⁺] = concentration of hydrogen ion and [A^-] = concentration of the conjucate base.
• Cross multiplying:

[H⁺][A^-] = K_a [HA]

[H⁺] = K_a * ( [HA] / [A^-])

• Taking -log of both sids

-log [H⁺] = -log  K_a -log  ( [HA] / [A^-] )

• Substitute pH and pK for -log[H⁺] and -logK respectively. The -log of the dissociation constant is defined as the pKa.

pH = pK_a  -log  ( [HA] / [A^-] ) (The stronger the acid ,lower is its pKa)

• Then to remove the -sign invert the last term

pH = pK_a  + log  ( [A^-] / [HA] )

• When the pH is equal to pKa, the concentration of conjugate base is equal to concentration of undissociated acid.