Chemistry: Solutions: Abnormal Colligative Properties, Degree of Association and Dissociation (For CBSE, ICSE, IAS, NET, NRA 2022)

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Abnormal Colligative Properties

  • The colligative properties of the solutions depend upon the number of solute particles present in the solution and not on their nature.
  • But sometimes while measuring colligative properties abnormal results are obtained due to the following reasons they are:
Abnormal Colligative Properties

Reason 1

If the solution is very concentrated, the particles of the solute interact with each other.

Reason 2

  • In case of association two or more solute molecules form a bigger molecule. The number of effective molecules in the solution decreases.
  • Consequently, the value of the colligative properties such as relative lowering of vapour pressure, elevation of boiling point, depression of freezing point, osmotic pressure are observed to be less than unassociated molecules.
  • The colligative property is inversely proportional to the molar mass, the molar mass of such solutes calculated on the basis of colligative property will be greater than the true molar mass of the solute.

Reason 3

  • In case of dissociation of the solute in the solution, the number of effective solute particles increases.
  • In such cases the value of the observed colligative property will be greater than that calculated on the basis of undissociated solute particles.
  • The molar mass of the solute calculated from the measurement of colligative property will be lower than the true molar mass of the solute.

Van ′ T Hoff Factor

  • In order to account for extent of association or dissociation the Van ′ t Hoff introduced a factor ′ i ′ .
  • i = Observed colligative property ÷ Normal (calculated or expected)

Colligative Property

  • The colligative property is proportional to the number of solute particles or the number of moles of solute i = Total number of moles of solute in the solution ÷ Expected (calculated) number of moles of solute
  • The colligative properties vary inversely as the molar mass of the solute, it follows that, i = Normal molar mass ÷ Observed molar mass
  • The observed molar mass is the experimentally determined molar mass whereas the normal molar mass is the molar mass calculated on the basis of chemical formula of the solute.
  • In case of association the value of Van ′ t Hoff factor is less than unity while dissociation is greater than unity.
  • The inclusion of Van ′ t Hoff factor, i, modifies the equations for the colligative properties as follows: (P0A – PA) ÷ P0A = I XB

ΔTb = i Kb m

ΔTf = i Kf m

πV = i CRT

Degree of Association

  • The degree of association may be defined as the fraction of the total number of molecules which associate to form a bigger molecule.
  • Let us consider the association of benzoic acid in benzene.
  • In benzene two molecules of benzoic acid associate to form a dimer.
  • It can be represented as 2C6H5COOH (C6H5COOH)2
  • If x represents the degree of association of benzoic acid in benzene (i.e.. out of one molecule of benzoic acid, x molecules associate to form a dimer) , then at equilibrium.
  • No. of moles of associated benzoic acid = 1 – x
  • No of moles of associated benzoic acid =
  • Total number of effective moles of benzoic acid =
  • According to definition, Van ′ t Hoff factor is given by i = Total number of moles of solute in the solution ÷ Expected (calculated) number of moles of solute =

Degree of Dissociation

  • The degree of dissociation may be defined as the fraction of the total number of particles that dissociate into simpler ions.
  • Consider a solution of KCl in water. When KCl is dissolved in water, it dissociates into K + and Cl ions. KCl K + + Cl
  • Let x be the degree of dissociation of KCl, then at equilibrium, number of moles of undissociated KCl = 1 – x
  • According to the dissociation of KCl shown above, when x mol of KCl dissociates, x moles of K + ions and x mol of Cl ions are produces
  • Thus, the total number of moles in the solution after dissociation = 1 – x + x + x = 1 + x
  • Hence, i = (Total number of moles of solute in the solution) ÷ (Expected (calculated) number of moles of solute) = .

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