Buffer Solutions: Buffer Action and Buffer Capacity

Buffer Solutions Defination, types, Mechanism of Buffer Action, Henderson's equation, Buffer Capacity and importance of buffer solutions with some math.

At ordinary temperature, pH of pure water is 7 (pH = 7). Now, when 1 mL of 1 (M) HCI solution is added to 100 mL of pure water, it is observed that the value of pH decreases from 7 to 2. Again, if 1 mL of 1 (M) NaOH solution is added to 100 mL of pure water, the value of the pH of the solution increases from 7 to 12. But, there are many solutions which resist the change of pH even when a small quantity of acid or base is added to them. Such solutions are called buffer solutions.

Definition of Buffer Solution: A buffer solution may be defined as the solution which resists the change in pH on addition of small amount of acid or base to it.


Various types of buffer solutions

[1] Solution of weak acid and its salt: Aqueous solution of a weak acid with its highly soluble salt (formed by reaction with a strong base) acts as a buffer solution.

Example 

  • Aqueous solution of CH3COOH and CH3COONa 
  • H2CO3 and NaHCO3
  • Citric acid and sodium citrate
  • Boric acid and sodium borate etc. are examples of buffer solutions.

[2] Solution of weak base and Its salt: Aqueous solution of a weak base with its highly soluble salt (formed by reaction with a strong acid) acts as a buffer solution.

Example

  •  Aqueous solution of NH3, with NH4Cl
  • Aniline and anilinium hydrochloride etc. are such type of Buffer solutions.

[3] Solution of two salts of a polybasic acid: Aqueous solutions of two salts of a polybasic acid can act as a buffer solution.

Example 

  • Aqueous solutions of Na2CO, and NaHCO3 (NaHCO3 is a weak acid and Na2CO3, is its salt) 
  • NaH2PO4 and Na2HPO4 etc.

Types of buffer solutions depending on their pH-values

Depending on the pH values of buffer solutions, they are divided into two classes, viz, acidic buffer and basic buffer.
Buffer Solutions

  1. Acidic buffer: The buffer solutions having pH value less than 7 are said to be acidic buffers. For example- aqueous solutions of CH3COOH and CH3COONa, aqueous solution of lactic acid and sodium lactate, etc. are some of familiar examples of acidic buffer solutions.
  2. Basic buffer: The buffer solutions having pH value more than 7 are called basic buffers. For example- an aqueous solution of NH4OH and NH4CI is a common example of basic buffer solution.

7.18.1 Mechanism of buffer action

Buffer action : The ability of a buffer solution to resist the change in pH on addition of small amounts of an acid or a base to it is called buffer action.

Mechanism of action of acidic buffer

Let us consider an acidic buffer consisting of a solution of weak acid, CH3COOH and its salt, CH3COONa. As the salt is a strong electrolyte, it dissociates almost completely

CH3COONa(aq) --------------> CH3COO-(aq) + Na+ (aq)

but CH3COOH being a weak acid, gets partially ionised in aqueous solution and exists in a state of equilibrium.

CH3COOH(aq) + H2O(l) = CH3COO-(aq)  + H3O+(aq)   ............[1]

Addition of a small amount of strong acid

If a small quantliy of strong acid (e.g., HCl, H2S04 etc.) is added to the buffer solutlon, H30+ ions derived from that strong acid combine with equal number of CH3COO- lons present in the solution to
form unionised molecules of CH,COOH:

CH3COO-(aq)  +  H3O+(aq) = CH3COOH(aq) + H2O(l)   ............(1)

As a result, the equilibrium involving ionisation of CH3COOH (equation 1) shifts towards left. Thus, there is practically no increase in the concentratlon of H30+ lons in the solutlon, i.e., pH of the solution remaIns almost unchanged.

Addition of a small amount of strong base

When a small quantity of a strong base like NaOH, KOH etc., is added to the buffer solution, the OH- ions originating from the strong base combine with equal number of H30+ lons present in the solution to produce unionised water molecules. Consequently, the equilibrium regurding the ionisation of CH3COOH [eqution (1)] gets disturbed. So, according to Le Chateller's principle, some molecules of CH3COOH ionise to produce H30+ ions and the equilibrium is re-established, Hence, addition of small quantity of strong base results netther any increase in the concentration of OH- ions nor any decrease In the concentratlon of H30+ ions, i.e, pH of the solution remains almost unaltered.

Mechanism of action of basic buffer

Let a basic buffer solution be composed of a weak base, NH3, and its salt, NH4Cl. Being a strong electrolyte, ammonium chloride (NH4CI) dissociates almost completely in solution:

NH4Cl(aq) ----------> NH4+(aq) + Cl-(aq).

But NH3, a weak base, reacts partially with water and forms the following
equilibrium (equation 2) in solution.

NH3 a weak base reacts partially with water and forms the following

Addition of a small amount of strong acid

If small quantity of strong acid (e.g., HCl, H2SO4 etc.) is added to that buffer solution, H30+ ions obtained from the strong acid react almost completely with equal number of OH- ions present in the
solution to produce unionised water molecules. As a result, the equiLibrium [eqn. (1)] gets disturbed. According to Le Chateller's princlple, some NH3 molecules react with water to form OH- Ions. As a result, the equilibrium [eqn. (1)] shifts towards right. So, due to addition of strong acid to the buffer solution, neither concentration of H+ ions increase nor the concentration of OH- ions decrease, i.e., pH of the solution remains unchanged.

Addition of a small amount of strong base

When a small quantity of a strong base (e.g., NaOH, KOH etc.) is added to the buffer solution, OH- ions produced by the strong base react almost completely with the equal number of NH4+ ions existing in the solution to form the unionised molecules of weak base, NH3 :

Buffer Action and Buffer Capacity


As a result, the equilibrium (equation. (1)] shifts towards left. Consequently, the concentration of OH- ions practically does not Increase, i.e, pH of the solution does not suffer from any change.

Buffer action of two salts of polybasic acid (in solution)

H3PO4 is a polybasic acid. In the buffer solution composed of two salts of this acid Na2H2PO4 and Na2HPO4, the former acts as a weak acid while the latter works as a salt. In the aqueous solution, these two salts undergo complete dissociation as follows :

NaH2PO4 (aq) →   Na+ (aq) + H2PO4- (aq)
Na2HPO4 (aq) → 2Na+ (aq) + HP042-  (aq)

H3PO4- itself is a weak acid. Moreover due to common ion [HPO4]2- effect, it undergoes slight ionisation to establish the following equilibrium:

H2PO4- (aq) + H2O(l) = [HPO4]2- (aq) + H30+(aq) .......[1]

Addition of a small amount of strong acid

When a small quantity of a strong acid is added to the buffer solutlon, the H3O+ ions produced by the strong acid react almost completely with an equal number of [HPO4]2-, ions to form unionised H2PO4-, ions as follows :

H2PO4- (aq) + H2O(l) <= [HPO4]2- (aq) + H30+(aq)

As a result, the above equilibrium (equation (1)] shifts towards left. Hence, the effect of the addition of a small amount of strong acid is neutralised. Thus, the pH of the solution remains undisturbed

Addition of a small amount of strong base

When a small quantity of a strong base is added to the buffer solution the OH- ions from the added base almost completely react with equal number of H3O+ ions present in solution to form unionised water molecules, causing disturbance to the ionisation equilibrium of NaH2PO4 [equation (1)]. Naturally, as expected from Le Chateller's principle, some molecules of NaH2PO4 ionise to form H3O+ ions and the equilibrium gets shifted towards right, Therefore, the additon of a small quantity of a strong base causes neither any increase in the concentration of OH- ions nor any decrease in the
concentration of H30+ ions, i.e., the pH of the solutions remains almost unchanged.

7.18.2 Determination of pH of a buffer solution: Henderson's equation

Let us consider a buffer solution composed of a weak acid (HA) and its salt (MA), In the solution, MA undergoes complete dissociation to form M+(aq) and A- (aq) ions and weak acid, HA partially ionises to establish the following equillibrium :

HA(aq) + H20(1) =  HO-(aq) + A+ (aq)  ..........(1)

Applying the law of mass action for the above equilibrium we get

 pH = pka+ log{[salt]/[acid]}      ...........(2)

where [acid] and [salt] are the initial molar concentrations of acid and salt respectively in the buffer solution.
Equation (2) is the Henderson's Equation to determine the pH of a buffer solution composed of a weak acid and its salt.

Similarly, the equation required to determine the pOH of a buffer solution consisting of weak base and its salt is- [salt]

pOH = pKb + log{[salt]/[base]}

where, Kb =the ionisation constant of weak base and [salt] and [base] are the respective initial molar concentrations of salt and base. In such cases,

pH = 14- pOH = 14- pkb- log{[salt]/[base]}
  • At a fixed temperature, the value of pK, of a weak acid is fixed. Hence, at a fixed temperature, pH of a buffer solution composed of a weak acid and Its salt depends upon the pKa of the weak acid as well as the ratio of molar concentration of the salt to the acid.
  • Similarly, at a particular temperature, pH of a buffer solution consisting of a weak base and its salt depends on the ratio of molar concentration of the salt to the base.

Application of Henderson's equation


  1. If the molar concentrations of a weak acid (or base) and its salt present in a buffer solution as well as the dissociation constant of that acid are known, pH of that solution can be determined by using Henderson's equation.
  2. If in an acidic or basic buffer solution, the molar concentrations of the weak acid (or base) and its salt and pH of that solution are known, then the dissociation constant of the weak acid can be calculated with the help of Henderson's equation.
  3. For the preparation of a buffer solution with a desired value of pH, the ratio between the weak acid and its salt (or weak base and its salt) in the solution can be determined by Henderson's equation.

7.18.3 Buffer capacity

Definition : Buffer capacity of a buffer solution is defined as the number of gram-moles of strong base or acid required to change the pH of 1 litre of that buffer solution by 1 unit.

Mathematical explanation: When the pH of 1 L of buffer solution increases by 'd(pH)' due to the addition of 'db' mol of any strong base, then the buffer capacity of that buffer solution,

(B) =(db)/d(pH)

Similarly, when the pH of 1L of buffer solution decreases by d(pH) due to the addition of db mol of any acid, then the buffer capacity of that buffer solution,
(B) =da/d(pH)

Some important facts regarding buffer capacity


  1. If the buffer capacity of any buffer solution is high, then greater amount of strong acid or base will be required to bring about any change in pH of that solution.
  2. Between two buffer solutions having same constituents, the buffer capacity of that one will be higher in which the concentrations of the constituents are higher.
  3. In a buffer solution, if the difference in molar concentration fo of the constituents is less, then addition of a small Concentrations or the constituents, Consequently, change when the molar concentrations of the constituents are equal to each other.

pH range of buffer capacity:

  1. A buffer solution consisting of a weak acid and its salt will work properly only when the ratio of the molar concentration of the salt to the acid lies in between 0.1 to 10. Therefore, according to Henderson's equation, pH-range of buffer capacity of such a buffer will vary from (pKa-1) to (pka +1)
  2. Similarly, range of pOH of buffer capacity for a buffer consisting of a weak base and its salt will vary from (pka - 1) to (pKa +1).

7.18.4 Importance of buffer solutions

[1] The value of pl of human blood lies in between 7,35 to 7.45, Le, human blood is slightly alkaline. In spite of the presence of acidic or basic substances produced by various metabolic processes, pH value of human blood remains fixed due to the presence of different buffer systems like bicarbonate-carbonic acid buffer (HCO3-/ H2CO3), phosphate buffer (HPO4 2- /H2PO4-) ete. The excess acid in the blood is neutralised by the reaction:  

H3PO3- +H3O+  = H2CO3

 H2CO3 decomposes into CO2 and water . The produced CO2 is exhaled through lungs. Again, if any base (OH-) from extraneous source enters blood, it gets neutralized by the reaction:

H2CO3 + OH- = HCO3- + H2O

[2] Buffer solutions find extensive use in analytical works in the laboratory as well as in chemical industries, In these cases, buffer solution is used to maintain the pH fixed at a certain value, For instance, in qualitative analysis in the laboratory, processes like electroplating, tanning of hides
and skins, fermentation, manufacture of paper, ink, paints and dyes, etc, pH is strictly maintained at a certain value. Buffer solution is also used for preservation of Fruits and fruit-products.

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Advanced Chemistry: Buffer Solutions: Buffer Action and Buffer Capacity
Buffer Solutions: Buffer Action and Buffer Capacity
Buffer Solutions Defination, types, Mechanism of Buffer Action, Henderson's equation, Buffer Capacity and importance of buffer solutions with some math.
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