Buffer Solutions

A solution whose pH is not altered to any great extent by the addition of small quantities of either acid (H⁺ ions) or a base (OH^– ions) is called the buffer solution. It can also be defined as a solution of reserve acidity or alkalinity which resists change of pH upon the addition of small amount of acid or alkali.

(1)     Types of buffer solutions : There are two types of buffer solutions,

(i)      Solutions of single substances : The solution of the salt of a weak acid and a weak base.

          Example : ammonium acetate (CH₃COONH₄), NH₄CN act as a buffer.

(ii)     Solutions of Mixtures : These are further of two types,

(a)     Acidic buffer : It is the solution of a mixture of a weak acid and a salt of this weak acid with a strong base.

          Example : CH₃COOH + CH₃COONa

(b)     Basic buffer : It is the solution of a mixture of a weak base and a salt of this weak base with a strong acid.

          Example : NH₄OH + NH₄Cl

(2)     Buffer action : Buffer action is the mechanism by which added H⁺ ions or OH^– ions are almost neutralised; so that pH practically remains constant. Reserved base of buffer neutralises the added H⁺ ions while the reserved acid of buffer neutralises the added OH^– ions.

The buffer action of different types of buffers may be explained as follows,

(i)      Buffer action of Ammonium acetate solution : Ammonium acetate like all other salts, is almost completely dissociated in the aqueous solution as follows,

CH₃COONH₄ → CH₃COO^– + NH₄⁺

Thus in the solution there is excess of CH­₃COO^– ions and NH₄⁺ ions.

When a few drops of an acid (say HCl) are added to the above solution, the H₃O⁺ ions given by the acid combine with the CH₃COO^–ions to form weakly ionized molecules of CH₃COOH.

CH₃COO^– + H₃O⁺ → CH₃COOH + H₂O

Thus the H₃O⁺ ion concentration of the solution does not change practically and hence the pH of the solution remains almost constant. Similarly, when a few drops of a base (say NaOH) are added to the above solution, the OH^– ions given by the base combine with the NH₄⁺ ions to form weakly ionized molecules of NH₄OH.

NH₄⁺ + OH^– → NH₄OH

Thus the OH^– ion concentration and hence the H₃O⁺ ion concentration or the pH of the solution remains almost constant.

(ii)     Buffer action of acidic buffer : In order to understand the buffer action, let us consider an acidic buffer such as a solution containing an equimolar amounts of acetic acid and sodium acetate. The solution contains a large number of sodium (Na⁺), acetate (CH₃COO^–) ions and also a large number of undissociated acetic acid molecules.

CH₃COOH(aq) ⇌ CH₃COO^–(aq) + H⁺(aq) ; CH₃COONa(aq) → CH₃COO^–(aq) + Na⁺(aq)

Suppose, a few drops of HCl are added to this buffer solution. This would provide hydrogen (H⁺) ions. These additional H⁺ ions would combine with the large reserve of CH₃COO^– ions to form undissociated acetic acid molecules.

                            

Since the additional H⁺ ions are neutralised by CH₃COO^– ions in the solution, there will be no change in its  value. The reserve basicity of the solution is due to acetate ions. On the other hand if a few drops of NaOH are added to the buffer solution, it would provide  ions. These OH^– ions will combine with H⁺=H ions present in the buffer solution to form unionised water molecules. this would result in the greater ionisation of acetic acid in order to restore the concentration of hydrogen ions to its original value.

                   CH₃COOH  ⇌ H⁺ + CH₃COO⁻ ;    

Since the additional OH⁻ ions are neutralised by acetic acid molecules, the pH of solution does not change appreciably. The reserve acidity of the solution is due to undissociated acetic acid molecules.

 

(iii)    Buffer action of basic buffer : In case of a basic buffer solution containing equimolar quantities of ammonium hydroxide and ammonium chloride, there is a large concentration of ammonium ions, chloride ions and undissociated ammonium hydroxide.

NH₄OH(aq) ⇌ NH₄⁺(aq) + OH^–(aq) ; NH₄Cl(aq) . NH₄⁺(aq) + Cl^–(aq)

When a few drops of HCl (aq) are added, the additional H⁺ ions are neutralised by OH^– ions present in the buffer.

                            

As some of OH^– ions from NH₄OH combine with H⁺ ions from the acid, it would result into greater ionisation of NH₄OH to restore the concentration of OH^– ions.

Thus, reserve basicity of the solution is due to the presence of undissociated ammonium hydroxide.

On the other hand, when a few drops of NaOH are added to the buffer solution, the additional OH^– ions are neutralised by ammonium ions.

Therefore, reserve acidity of solution is due to the presence of ammonium ions.

 Basic buffer (NH₄OH + NH₄Cl)

Diagrammatic representation of buffer action of Basic buffer

 

(3)     Examples of buffer solutions

(i)      Phthalic acid + potassium hydrogen phthalate

(ii)     Citric acid + sodium citrate.

(iii)    Boric acid + borax (sodium tetraborate).

(iv)    Carbonic acid (H₂CO₃)⁺  sodium hydrogen carbonate (NaHCO₃). This system is found in blood and helps in maintaining pH of the blood close to 7.4 (pH value of human blood lies between 7.36 – 7.42; a change in pH by 0.2 units may cause death).

(v)     NaH₂PO₄ + Na₃PO₄

(vi)    NaH₂PO₄ + Na₂HPO₄

(vii)   Glycerine + HCl

(viii)  The pH value of gastric juice is maintained between 1.6 and 1.7 due to buffer system.

(4)     Henderson – Hasselbalch equation : pH of an acidic or a basic buffer can be calaculated by Henderson- Hasselbalch equation.

For acidic buffers,    

When , then,  pH = 1 + pK_a and when , then,  pH = pK_a – 1

So weak acid may be used for preparing buffer solutions having  values lying within the ranges pK_a + 1 and pK_a – 1. The acetic acid has a pK_a of about 4.8; it may, therefore, be used for making buffer solutions with pH values lying roughly within the range 3.8 to 5.8.

For basic Buffers,

Knowing pOH, pH can be calculated by the application of formula, pH + pOH = 14

pH of a buffer solution does not change with dilution but it varies with temperature because value of K_w changes with temperature.

(5)     Buffer capacity : The property of a buffer solution to resist alteration in its pH value is known as buffer capacity. It has been found that if the ratio or is unity, the pH of a particular buffer does not change at all. Buffer capacity is defined quantitatively as number of moles of acid or base added in one litre of solution as to change the pH by unity, i.e.,

Buffer capacity

  or 

Where  is number of moles of acid or base added to 1 litre solution and ΔpH is change in pH.

Thus greater the buffer capacity, the greater is its capacity to resist change in pH value. Buffer capacity is greatest when the concentration of salt and weak acid/base are equal, or when pH = pK_a or pOH = pK_b.

(6)     Significance of buffer solutions

(i)      Buffer solutions are used for comparing colorimetrically the hydrogen ion concentration of unknown solutions.

(ii)     Acetic acid-sodium acetate is used in the removal of phosphate radical during the qualitative analysis of the mixture.

(iii)    NH₄Cl/NH₄OH buffer is used for the precipitation of hydroxides of third group of qualitative analysis.

(iv)    In industries, buffer solutions are used in the alcoholic fermentation (pH 5 to 6.5), tanning of leather, electroplating, manufacture of sugar, paper manufacturing etc.

(v)     In bacteriological research culture media are generally buffered to maintain the pH required for the growth of the bacteria being studied.

(vi)    In biological systems buffer system of carbonic acid and sodium bicarbonate is found in our blood. It maintains the pH of blood to a constant value (about 7.4) inspite of various acid and base-producing reactions going on in our body.