Henderson – hasselbalch equation The Henderson – Hasselbalch equation describes the derivation of pH as a measure of acidity ( using pKa , the acid dissociation constant ) in biological and chemical systems. The equation is also useful fro estimating the pH of a buffer solution and finding the equilibrium pH in acid - base reaction ( it is widely used to calculate isoelectric point of the proteins)

limitations The most significant is the assumption that the concentration of the acid and its conjugate base at equilibrium will remain the same as formal concentration . This neglects the dissociation of the acid and the hydrolysis of the base . The dissociation of water itself is neglected as well . These appropriate will fail when relatively strong acid or bases(pKa more a couple units away from 7),dilute or very concentrated solution ( less than 1 mM or greater) , or heavily skewed acid /base ration ( more than 100 to 1)

Steps in the Preparation of Buffers Determine the pH range required to create the buffer.  For example, a laboratory test needs to take place at pH 4.00. So you have to create a buffer that will remain at pH 4.00 despite the addition of acid or base solutions.

Use the Henderson-Hasselbach equation and substitute the pKa of buffer and pH needed. In my example, pH = 4.00 and pka= 3.75

Check the table of Weak acids and their conjugate base for the buffer that will perform best on the required pH range. In choosing the buffer to use, the pKa of the buffer should have the nearest value to your required pH.  In the given example, after looking at the table, the best buffer in pH 4.00 is formic acid-formate mixture because its pKa is equal to 3.75 .

Finding the concentration of the acid/base ratio is a two-step process: (4a) Calculate the difference between pH and pKa.  Applying this to the example, pH = 4.00 and pKa = 3.75 would give 4.00 ; 3.75 = 1.78  .Then calculate the antilogarithm in the Henderson-Hasselbach equation on step 2. Use your calculator to find the antilogarithm value

Prepare the chemical reagents that correspond to the concentrations of your base/acid ratio. Mix them in an equal amount or ratio depending on the volume you need to use. In my example, 25 mL of 0.1M of Formic acid and 25 mL of 0.0178M of sodium formate will yield a total of 50 mL buffer solution.

A buffer works by removing hydrogen ions from the solution when they are in excess and donating hydrogen ions to the solution when they are low. Most buffers are weak acids and their salts which dissociate in solution as shown in the following example: