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>> Relative Strength of Acids and Bases-Part1
Relative Strength of Acids and Bases:
According to Arrhenius concept, an acid is a substance which furnishes H+ ions when dissolved in water. All the acid properties on an acid are due to H+ ions present in the solution.
The extent to which an acid property is given by an acid is a measure of its acid strength. The acid strength of a solution does not depend on its concentration but on the number of H+ ions present. The concentration of H+ ions depends on the ionisation of an acid in solution. On dilution, the ionisation increases and more of H+ions come to solution with the result that the acid strength increases. Thus, acid strength increases on dilution while its concentration decreases.
At infinite dilution the dissociation of an acid is nearly complete and all acids are equally strong at infinite dilution.
The concentration of H+ ions at all other dilutions of equimolar solutions of the acids may not be equal and depends on their degree of dissociation. Thus, to measure the relative acid strength of the two acids, the measurements of hydrogen ion concentration, i.e., degree of dissociation is made of equinormal solutions of the two acids. Various methods are used for this purpose. Some are described below.
(i) The conductivity method: The degree of dissociation of a weak acid is equal to conductivity ratio Λ1/Λ∞. Thus, the degrees of dissociation a, and a2 for two equinormal acids are given by:
For acid I, α1 = Λ1/Λ∞
For acid II, α2 = Λ2/Λ∞2
At infinite dilution, all weak electrolytes have almost the same value of Λ∞; hence,
Λ∞1 = Λ∞2
(Strength of acid I)/(Strength of acid II)=α1/α2 = Λ1/Λ2 = ((1000×sp.cond.acid I)/C)/((1000×sp.cond.acid II)/C)
= (Sp.cond.acid I)/(Sp.cond.acid II)
The relative strength of two acids is, thus, equal to the ratio of their equivalent conductance or specific conductance of equinormal solutions which can be determined experimentally.
(ii) Comparing dissociation constants: Let K1 and K2 be the dissociation constants of two acids and let and a2 be their degree of dissociation in equinormal solutions.
Applying Ostwald's dilution law, α1 = √K1/c and α2 = √K2/c
Thus, (Strength of acid I)/(Strength of acid II) = α1/α2 = √(K1/K2 )
Dissociation constants of some weak acids are given in the table:
Acid-ionisation Constants at 25oC
Hydrogen Sulfate ion
(iii) Thomson thermal method: In this method, heat of neutralisation of two acids is first determined separately with NaOH. Let it be 'x' and 'y' calorie. The one gram equivalent of each of the two acids is mixed and one gram equivalent of NaOH is added. Let the heat evolved in this case be 'z' calorie. The two acids will neutralise a fraction of the base proportional to their relative strength. Suppose n gram equivalent of NaOH is neutralised by acid I and the rest (1 - n) by acid II.
Total Heat solved, z = nx + (1-n)y
= nx + y - ny
or z-y = n(x-y)
or n = (z-y)/(x-y)
So (Strength of acid I)/(Strength of acid II) = n/((1-n) ) = (((z-y))/((x-y) ))/(1-((z-y))/((x-y) )) = ((z-y))/((x-z) )
Relative strength of bases: A base is a substance which gives OH" ions when dissolved in water. The base strength depends on OH- ion concentration. The above methods can be used for measuring relative base strengths also. In the Thomson thermal method, the two bases and their mixtures will be neutralised by strong acid, say HC1.
The relative strengths of some of the acids are as follows:
(i) HCIO4 > HBr > HC1 > HNO3 > H2SO4 > H3O+ > H2SO3 > H2CO3>CH3COOH
(ii) HCIO4 > HCl03 > HClO2 > HCIO
(iii) HI > HBr > HCl > HF
(iv) HClO3 > HBrO3 > HIO3
(v) CCI3COOH > CHCl2COOH > CH2C1COOH > CH3COOH
(vi) HCOOH > CH3COOH > C2H5COOH
The relative strengths of some of the bases are as follows:
(i) KOH > NaOH > Ca(OH)2 > NH4OH
(ii) (CH3)2NH > CH3NH2 > (CH3)3N > NH3
(iii) (C2H5)2NH > C2H5NH2 > NH3 > (C2H5)3N
(v) NH3 > NH2NH2 > NH2OH
(vi) NH3 > C5H5N > C6H5NH2