When a solute is dissolved in a solvent, the freezing temperature is lowered in proportion to the number of moles of solute added. This property, known as freezing-point depression, is a colligative property; that is, it depends on the ratio of solute and solvent particles, not on the nature of the substance itself. The equation that shows this relationship is:

\Delta T = {K_f} \cdot m

where ΔT is the freezing point depression, Kf is the freezing point depression constant for a particular solvent (3.9°C-kg/mol for lauric acid in this experiment1 ), and m is the molality of the solution (in mol solute/kg solvent).

In this experiment, you will first find the freezing temperature of the pure solvent, lauric acid, CH3(CH2)10COOH. You will then add a known mass of benzoic acid solute, C6H5COOH, to a known mass of lauric acid, and determine the lowering of the freezing temperature of the solution. In an earlier experiment, you observed the effect on the cooling behavior at the freezing point of adding a solute to a pure substance. By measuring the freezing point depression, ΔT, and the mass of benzoic acid, you can use the formula above to find the molecular weight of the benzoic acid solute, in g/mol.


In this experiment, you will

  • Determine the freezing temperature of pure lauric acid.
  • Determine the freezing temperature of a solution of benzoic acid and lauric acid.
  • Examine the freezing curves for each.
  • Calculate the experimental molecular weight of benzoic acid.
  • Compare it to the accepted molecular weight for benzoic acid.