why do u think elevation in b.p. and depression in f.p. is directly prop. to molality but osmotic pressure to molarity?

also explain what is clausius-clapeyron eqn. plz. reply faster

F.P.  depression describes the phenomenon that the freezing point of a liquid (a solvent) is depressed when another compound is added, meaning that a solution has a lower freezing point than a pure solvent.

This happens whenever a solute is added to a pure solvent, such as water. The phenomenon may be observed in sea water, which due to its salt content remains liquid at temperatures below 0°C, the freezing point of pure water.

The freezing point depression is a colligative property, which means that it is dependent on the presence of dissolved particles and their number, but not their identity. It is an effect of the dilution of the solvent in the presence of a solute. It is a phenomenon that happens for all solutes in all solutions, even in ideal solutions, and does not depend on any specific solute-solvent interactions.

The freezing point depression happens both when the solute is an electrolyte, such as various salts, and a nonelectrolyte. In thermodynamic terms, the origin of the freezing point depression is entropic and is most easily explained in terms of the chemical potential of the solvent.

The extent of freezing-point depression can be calculated by applying Clausius-Clapeyron relation and Raoult's law together with the assumption of the non-solubility of the solute in the solid solvent. The result is that in dilute ideal solutions, the extent of freezing-point depression is directly proportional to the molal concentration of the solution according to the equation

ΔT_f = K_f · m_B

where

• ΔT_f, the freezing point depression, is defined as T_{f (pure solvent)} − T_{f (solution)}, the difference between the freezing point of the pure solvent and the solution. It is defined to assume positive values when the freezing point depression takes place.


• K_f, the cryoscopic constant, which is dependent on the properties of the solvent. It can be calculated as K_f = RT_m²M/ΔH_f, where R is the gas constant, T_m is the melting point of the pure solvent in kelvin, M is the molar mass of the solvent, and ΔH_f is the heat of fusion per mole of the solvent.


• m_B is the molality of the solution, calculated by taking dissociation into account since the freezing point depression is a colligative property, dependent on the number of particles in solution. This is most easily done by using the van 't Hoff factor i as m_B = m_solute · i. The factor i accounts for the number of individual particles (typically ions) formed by a compound in solution. Examples:
  
  
  
  • i = 1 for sugar in water
  
  
  • i = 2 for sodium chloride in water, due to dissociation of NaCl into Na⁺ and Cl^-
  
  
  • i = 3 for calcium chloride in water, due to dissociation of CaCl₂ into Ca²⁺ and 2 Cl^-
  
  
  • i = 2 for hydrogen chloride in water, due to complete dissociation of HCl into H⁺ and Cl^-
  
  
  • i = 1 for hydrogen chloride in benzene, due to no dissociation of HCl in a non-polar solvent
  
  
  
  

At high concentrations, the above formula is less precise due to the approximations used in its derivation and any nonideality of the solution. If the solute is highly soluble in the solid solvent, one of the key assumptions used in deriving the formula is not true. In this case the effect of the solute on the freezing point must be determined from the phase diagram of the mixture.

http://en.wikipedia.org/wiki/Clausius-Clapeyron_relation