Coupon Accepted Successfully!


Ionic Equilibria and Acids and Bases

Equilibrium involving ions are known as Ionic Equilibria. Acids, bases and salts come under this category.
So, the next question that comes spontaneously to our mind is, how do we define acids and bases.
The first theory was given by Arrhenius.

According to him, acids are substances that donate a proton, H+ in water.
Thus, HCl, H2SO4, HNO3, CH3COOH, all fall under acids according to this theory.
According to Arrhenius, bases are substances that donate a Hydroxyl Ion in water.
Hence, NaOH, Mg(OH)2, Al(OH)3 are bases according to Arrhenius.

Cannot explain acidic and basic nature in non aqueous medium
Cannot explain the basic character of NH3 and several molecules that do not possess OH group
Similarly cannot explain the acidic nature of B, Al compounds

Bronsted Lowry theory of Acids and Bases
Acids: capable of donating a proton to give a conjugate base
Bases: capable of accepting a proton to give a conjugate acid
Explains, the basic nature of NH3
According to this theory, the conjugate base of a weak acid is strong and vice versa
Similarly, the conjugate acid of a strong base is weak and vice versa
This concept explains the strength of acids and bases based on how readily they accept or donate H+
Limitation cannot explain the acidic nature of B, Al compounds

Lewis Acids and Bases
According to this theory,
Acids are electron deficient compounds that accept a lone pair of electrons, and bases are compounds that donate a pair of electrons to form adducts.
Thus, explains the acidic nature of B, Al compounds and basic nature of electron rich compounds such as NH3

Ionization of Acids and Bases
An Acid HA ionizes in water as
HA + H2O → H3O+ + A-
Where, A- is the conjugate base of the acid HA.

Ionization constant of Water
H2O + H2O H3O+ + OH-
The dissociation constant, Kw = ([H3O+][OH-])/[H2O]2
However, since the [H2O] does not change for all practical reasons, as self ionization of water is very low, we omit the term, [H2O]2 in the denominator
This ionization is dependent on temperature and at 298 K, the value of Kw is found to be 1.00 ×10-14 M2
Thus, water largely exists in unionized form
pH: pH is defined as negative logarithm of the concentration of H+ ion
Mathematically, pH = - log [H3O+]

Ionization constants of weak acids
For weak acids,
HX + H2O H3O+ + X-
Let the initial concentrations of HX be a
The initial concentration of H3O+ and X- will respectively be 0

If, α is the extent of ionization, then; the concentration of HX at equilibrium will be, a (1-
α) and the concentrations of H3O+ and X- will be aα
Thus, Ka = [H3O+][ X-]/[HX]
= (a
Since for weak acids, α <<1, the above equation reduces to, Ka = aα2/(1-α)

Larger, the value of Ka, larger is the strength of the acid
We can write a similar expression for weak bases as well, where Ka is replaced by Kb

Dissociation constant of base
Ka and Kb for a given species are related as
Ka×Kb = Kw

Factors affecting the strength of an Acid
  1. Strength of H-A bond, the weaker the bond, the greater is the acid strength.
  2. Nature of Polarity. The greater the electronegativity difference between H and A, the larger the acid strength.
    However, within the same group, say Halogens,H-X bond strength overweighs the H-X Polarity.
    Thus, HI is the strongest acid and HF is the weakest.
Common Ion effect
When a strong acid or base is added to a weak acid or base, the degree of ionization of the weak acid or base is further suppressed.

For example, consider the equation,

When, a strong acid such as HCl is added to the above solution, then, Ac- being a very strong conjugate base of weak Acid HAc will readily accept H+ ions donated from the Strong acid HCl and therefore the equilibrium will shift towards the left, the undissociated form.

Hydrolysis of Salts
Salts are formed by combination of acids and bases

Case I: Salts of strong acids and strong bases - the pH of these solutions will be 7.0
Examples, HCl + NaOH NaCl + H2O
These salts do not ionize in water and therefore, their pH remains 7.0

Case II: Salts of Strong acids and Weak bases
The pH of these solutions will be Acidic, for example- HCl + NH4 OH NH4Cl + H2O
NH4Cl + H2O NH4+ + Cl-
The NH4+ ions combine with water to give,
NH4+ + H2O NH4OH + H+

Since, NH4OH is a weak base, it largely exists in unionized form and the concentration of H+ Ions increases rapidly. Thus the pH is always < 7.0.

Salts of weak acids and strong bases
Consider CH3COONa Ac- and Na+ in aqueous solutions.
Ac- + H2O HAc + OH-
Since, HAc is a weak acid, it largely remains in undissociated form, and thus [OH-] increases rapidly. Hence, the solutions of these salts have a pH >7.0.

For salts of weak acids and weak bases
pH = 7.0 + ½(pKa + pKb) and the pH of these solutions depends on the relative strengths of the respective weak acids and weak bases
Buffer solutions resist a change in pH on addition of a small amount of acids and bases.
Solubility Product Constant is given by Ksp and is determined for sparingly soluble salts.
If the solubility product is greater than ionic product, the substance is said to be soluble in water and if the solubility product is lesser than ionic product of water, then, the substance would precipitate in water.

Test Your Skills Now!
Take a Quiz now
Reviewer Name