Literature Review 3: Investigation on how dissolving chemicals in water changes the freezing point of a solution.

Project Title 3: Investigation on how dissolving chemicals in water changes the freezing point of a solution.

Hypothesis:  The higher amount of dissolving chemicals in water, the lower the freezing point of water.

Done by: Liew Yee Theng

Through some research, salt, otherwise known as sodium chloride (CaCl2), lowers the freezing point of water. This process is known as freezing point depression and it can occur as long as we have a solution. Adding a solute, like salt, to a solvent, like water, lowers the freezing point of the solvent. The answer to how much the freezing point of the solution depends on three things: the molality of the solution, the van't Hoff factor of the solute, and the molal freezing-point-depression constant of the solvent. 

Molality, m, is defined as moles (mol) solute per kilograms (kg) solvent, as shown in Equation 1, below: 

Equation 1:

Molality (moles/kg) = Moles (mol) of Solute/Kilograms (kg) of Solvent

Freezing point depression is a colligative property, a property that depend on how many solute particles are in the solvent, not the kind of solute particles. Molality, m, is one piece of this "how many solute particles are present?" question. The van't Hoff factor is the second part of the "how many solute particles are present?" question.

The van't Hoff factor, i, deals with how a molecule of solute dissociates, or breaks apart, in the solvent. Covalent compounds, like sucrose (C₁₂H₂₂O₁₁), do not dissociate in solution. These compounds have van't Hoff factors i = 1. Ionic compounds, like table salt (NaCl), dissociate when in solution. Table salt (NaCl) has a van't Hoff factor i = 2 because it dissociates into two ions in solution, Na⁺ and Cl^-.

The third factor, the molal freezing-point-depression constant, K_f, is different for every solvent. It has units of (° C/m), and it tells us how much 1 mol of solute added to 1 kg of solvent will lower the solvent's freezing point. For pure water, K_f = 1.86° C/m.

Combining these three factors—molality, m, van't Hoff factor, i, and molal freezing-point-depression constant, K_f—into an equation that predicts how much the freezing point of a solvent will decrease, ΔT, when a certain amount of solute is added. Equation 2, below, is the freezing point depression equation:

Equation 2:

Degrees Freezing Point is Depressed (° C) = Molal Freezing-Point-Depression Constant (° C/m)
× molality of solution (mol solute/kg solvent) × van't Hoff Factor (unitless)

ΔT = K_f m i

ΔT is the freezing point depression in degrees Celsius (° C) Kf is the molal feezing-point-depression constant in degrees Celsius per molal (° C/m) m is the molality of the solution in moles per kilogram (mol/kg) i is the van't Hoff factor of the solute, which does not have units

In order to find out the new freezing point of a solution, T_n, subtract the change in temperature, ΔT, from the original freezing point, T_f, as shown in Equation 3, the solution freezing point equation:

Equation 3:

Solution Freezing Point (° C) = Solvent Freezing Point (° C) - Degrees Freezing Point is Depressed (° C)
T_n = T_f - ΔT

Tn is the freezing point of the solution in degrees Celsius (° C) Tf is the freezing point of the solvent in degrees Celsius (° C) ΔT is the freezing point depression in degrees Celsius (° C)

Bibliography

• Eli, Todd & Keith. (n.d.). "Colligative Properties," Chemworld, ThinkQuest Library, Oracle Education Foundation. Retrieved August 13, 2012, fromhttp://library.thinkquest.org/C006669/data/Chem/colligative/colligative.html?tqskip1=1.

• Lachish, U. (2000). Avogadro's Number, Atomic and Molecular Weight. Retrieved August 13, 2012, from http://urila.tripod.com/mole.htm.

This idea for this Project Idea came from this source.

• MacQuade, J., et al. (1986). It's Getting Colder (Freezing Point Depression). Retrieved August 13, 2012, fromhttp://www.woodrow.org/teachers/chemistry/institutes/1986/exp9.html.