CHAPTER 6

IONIC EQUILIBRIUM

6.1      Concepts of Acid and Base

6.1.1 Arrhenius concept

Acid is a substance which is capable of furnishing H+ ions in aqueous solution. e.g.-HCl, H2SO4 etc. While base is a substance which furnishes OH– ions. e.g. - NaOH, NH4OH etc.

Limitations: Arrhenius theory failed to explain.

(i)      Behaviour of acids/bases in non aqueous solutions.

(ii)     Neutralization reaction giving salt in absence of a solvent.

(iii)    Acid character of certain salts like AlCl3,BF3 etc.

(iv)    Existence of H+ in water.

 

6.1.2 Bronsted Lowery Concept

Acid is a substance which is capable of donating a proton while base is a substance which is capable of accepting a proton. This is also called protonic theory of acids-bases.

HCl        H2O       ®           H3O+              +            Cl–

acid        base                conjugate acid                 conjugate base

H2O  +  NH3        ®                           +            OH–

acid        base                conjugate acid                 conjugate base     

conjugate base of an acid is a species formed by the loss of a proton from acid.

                              Acid ® H⁺ + conjugated base

Similarly conjugate acid is formed from a base by gain of H+.

                              Base + H⁺ ® conjugated acid

Weak acid has a strong conjugate base and vice-versa.

A Bronsted - Lowery acid base reaction always proceeds in the direction from the stronger to the weaker acid base combination. e.g.

HI          +             OH–       ®           H2O                 +          I–

Strong                   Strong                  Weak                           Weak

acid                       base                      acid                              base

 

6.1.3 Lewis concept

Acid is a substance which can accept a pair of electrons while base is a substance which can donate a pair of electrons.

Hence Lewis acids are

(i)      Molecules in which central atom has incomplete octet e.g. BF3, AlCl3 and FeCl3 etc.

(ii)           Simple cations like Ag+, H+ etc.

(iii)    Molecules in which central atom has vacant d-orbitals e.g.- SiF4, SnCl4 etc. 

(iv)    Molecules in which atoms of different electronegativities are joined by multiple bond. e.g. CO2, SO3 etc.

(v)           In carbonyl complexes, metal atoms act as Lew’s acids e.g. Ni in Ni(CO)4 etc.

And Lewis bases are

(i)      Neutral molecules like NH3, RNH2 etc.

(ii)           Negatively charged anions. e.g. CN–, Cl– etc.

(iii)    Molecules with carbon-carbon multiple bonds can act as lewis base in some cases - e.g.  in

(iv)    In complex compounds, the ligands act as Lewis bases e.g. CO in Ni(CO)4 etc.

 

6.2      Strong and weak electrolytes

Electrolytes which dissociate completely into ions in aqueous solutions are called strong electrolytes. Eg: HCl, H2SO4, NaCl etc. Whereas electrolytes are those which dissociate to a lesser extent are called weak electrolytes. Eg: CH3COOH, NH4OH etc.

 

6.3      Ionization of Weak Acids

Let us consider the ionization of weak acid HA, having initial concentration ‘c’ in mol/litre

                

It K_a is ionization constant for acid, then

         

              

for very weak acids. x < < 1, concentration can be calculated from following formula:

 (i)    

(ii)    

(iii)  

(iv)   

 

6.4 Common Ion Effect

In presence of common ion dissociation of weak electrolyte is further suppressed. e.g.

CH3COOH is weak electrolyte. It dissociates as

If some amount of (CH₃COONa) is added to the solution of acetic acid. Then  will provide common acetate ions. Consequently equilibrium will shift towards reactant side and degree of dissociation becomes further smaller.

 

6.5 Mixture of two weak acids

Let us consider the mixture of aqueous solutions of two weak acids HA₁ and HA₂ whose concentrations and are C₁ and C₂ and ionization constants are Ka₁ and Ka₂.

      HA₁                       H⁺        +          A₁^—

          C₁                                   —                    —

          C₁(1 – x₁)                                   C₁x₁                 C₁x₁                 (At equilibrium)

          HA₂                       H⁺        +          A₂^—

          C₂                                   —                    —

          C₂(1 – x₂)                                   C₂x₂                 C₂x₂                             (At equilibrium)

                                              …(i)

and                                             …(ii)

Dividing Eq. 1 by Eq. 2

  on putting the value of x₁ in terms of x₂ in equation (i), we can calculate x₂ from above discussion we can calculate the pH of solution mixture of two acids.                                

 

6.6      Ionization of polyprotic acids

Let us consider the ionization of H₂S in its aqueous solution. It C is concentration in mole/litre C and Ka₁ and Ka₂ are ionization constants for first and second step ionization of H₂S.

             H₂S                    H⁺                    + HS^—

             C(1-x₁)           cx₁                   cx₁(1 – y₁)        (At equilibrium)

             HS         H⁺          +          S^{– –}

       cx₁ (1–y₁)           cx₁ y₁                  x₁ y₁                 (At equilibrium)

                                          

                                       

because generally  hence we can assume that  

\                                                   …(i)

                                                      …(ii)

From equation (i) and (ii) we can calculate the pH of aqueous solution of H₂S.

 

6.7      Ionisation of Water

Ionic product of water ()

Water is weak electrolyte, hence

        

        

        

Since [H2O] = constant =

 depends on temperature. At 25° C value of   is 10–14.

Hence

Thus

Or   

Since [H+] = [OH–]

At 25°C, PH of pure water 

Now at 25°C

or    

or  at 25°C.

Therefore PH range at 25° C will be 0 to 14

As temperature increases, degree of dissociation of water also increases, therefore value of   increases.

 

6.8 Buffer solutions

The solutions which resist the change in its pH value on addition of small amount of acid or base are called buffer solutions. On adding small amount of acid or base there is no significant Change in pH of the buffer solution.

 

(a) Acidic Buffer:

It contains mixture of weak acid and its salt with strong base. e.g. mixture of CH3 COOH and CH3COONa.

         For acidic buffer of

         We have in solution -  molecules, CH3COO– ions and Na+ ions.

 Dissociation of weak electrolyte is suppressed in the presence of common ion CH₃COO^– from CH₃COONa. So, p^H  for such buffer can be calculated by following formula:

              

 

(b)           Basic Buffer:        

It is the mixture of weak base and its salt with strong acid. e.g. NH₄OH and NH₄Cl.

For basic buffer solution of NH₄OH and NH₄Cl. We have in solution- NH₄OH,  ions and Cl– ions.

Dissociation of weak electrolyte NH₄OH is suppressed in the presence of common ion  from NH₄Cl. So, p^OH  for such buffer can be given by:

 

          p^OH =                 

 

6.9 Solubility Product

It is the product of the molar concentrations of the ions in a saturated solution of an sparingly soluble salt with each concentration term raised to the power equal to the number of times that ion appears in balanced equation that represents equilibrium. It is denoted by Ksp.

           

         

In saturated solution [A_xB_y] = constant

         

         

Applications of Solubility Product

(i)      It helps in predicting the formation of a precipitate.

               If ionic product > Ksp, precipitation occurs and if ionic product < Ksp no precipitate is formed.

(ii)     Calculation of solubility of sparingly soluble salt - let solubility of salt AxBy in water at a particular temperature is S mole per litre. Then at equilibrium.

               A_xB_y  

              

e.g. for   AgCl  

               Ksp = S2

Hence   

For         Ag₂CrO₄  

              

              

 (iii)   In qualitative analysis

          The separation and identification of various basic radicals into different groups is based upon (a) solubility product principle and (b) common ion effect.

(iv)    Purification of common salt.

          Saturated solution of impure common salt is prepared and insoluble impurities are filtered off. HCl gas is passed through this solution. Thus, ionic product of  exceeds the Ksp and pure NaCl precipitates out from the solution. This process is called salting out.