Boiling point of a solution is always higher than that of the pure solvent. For dilute solutions, elevation of boiling point is directly proportional to the molar concentration of solute in a solution.
Here, m is molality and Kb is called Boiling Point Elevation Constant or Molal Elevation Constant (Ebullioscopic Constant). The unit of Kb is K kg mol-1
If w2 g of solute of molar mass M2 is dissolved in w1 g of solvent. Then molality of solution is given as follows:
Substituting the value of molality in previous equation we get:
Example: 18 g of glucose C6H12O6 is dissolved in 1 kg of water in a saucepan. At what temperature will water boil at 1.013 bar? Kb for water is 0.52 K kg mol-1.
Answer: Moles of glucose `=(18g)/(180g\text(mol)^(-1))=0.1` mol
Mass of solvent = 1 kg
Hence, molality of glucose solutin = 0.1 mol kg-1
For water, change in boiling point
Boiling point of water at 1.013 bar is 373.15 K
So, boiling point of solution `= 373.15 + 0.052 = 373.202` K
Example: The boiling point of benzene is 353.23 K. When 1.80 g of a non-volatile solute was dissolved in 90 g of benzene, the boiling point is raised to 354.11 K. Calculate the molar mass of the solute. Kb for benzene is 2.53 K kg mol-1.
Answer: ΔTb = 354.11 K – 353.23 K = 0.88 K
So, molar mass of solute can be calculated as follows:
`=(1000xx1.8xx2.53)/(0.88xx90)=57.5` g mol-1
When a non-volatile solid is added to the solvent, its vapor pressure decreases and becomes equal to that of solid solvent at lower temperature. This results in lowering of freezing point of the solvent.
Depression of freezing point for dilute solute is directly proportional to molality.
Or, `ΔT_t=K_t\m` ………(1)
Here, the constant of proportionality, K1, is called Freezing Point Depression Constant, Depression Constnat or Cryoscopic Constant. The unit of Kf is K kg mol-1
We know that molality is given by following equation:
So, the equation for ΔTf can be written as follows:
`ΔT_f=(K_f\xx\w_2xx1000)/(ΔT_f\xx\w_1)` ……… (2)
Example: 45 g of ethylene glucol (C2H6O2 is mixed with 600 g of water. Calculate (a)the freezing point depression and (b) the freezing point of the solution.
Answer: First of all, we need to find molality of ethylene glycol.
Moles of ethylene glycol `=(45g)/(62g\text(mol)^-1)=0.73` mol
Mass of water `=(600g)/(1000kg)=0.6` kg
So, molality of ethylene glycol `=(0.73text(mol))/(0.60 kg)=1.2` mol kg-1
Now, change in freezing point can be calculated as follows:
`ΔT_f=1.86kg\text(mol)^(-1)xx1.2 text(mol)kg^(-1)=2.2` K
So, freezing point of aqueous solution `= 273.15 K – 2.2 K = 270.95` K
Example: 1.00 g of a non-electrolyte solute dissolved in 50 g of benzene lowered the freezing point of benzene by 0.40 K. The freezing point depression constant of benzene is 5.12 K kg mol-1. Find the molar mass of the solute.
>Answer: Molar mass of solute can be calculated as follows:
`M_2=(5.12 kg\text(mol)^(-1)xx1.00g\xx1000\g\kg^(-1))/(0.40xx50g)=256` g mol-1
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