🤓 Based on our data, we think this question is relevant for Professor Weber's class at UNT.

How many grams of ethylene glycol, C _{2}H_{4}(OH)_{2}, must be added to 400.0 g of water to yield a solution that will freeze at -8.35 °C?

a) 37 g

b) 111 g

c) 75 g

d) 151 g

We’re being asked to **determine the mass of ethylene glycol (C _{2}H_{4}(OH)_{2})** that must be added to

Recall that the freezing point of a solution is *lower* than that of the pure solvent and the ** change in freezing point (ΔT_{f})** is given by:

$\overline{){{\mathbf{\Delta T}}}_{{\mathbf{f}}}{\mathbf{=}}{{\mathbf{T}}}_{\mathbf{f}\mathbf{,}\mathbf{}\mathbf{pure}\mathbf{}\mathbf{solvent}}{\mathbf{-}}{{\mathbf{T}}}_{\mathbf{f}\mathbf{,}\mathbf{}\mathbf{solution}}}$

The ** change in freezing point** is also related to the molality of the solution:

$\overline{){{\mathbf{\Delta T}}}_{{\mathbf{f}}}{\mathbf{=}}{{\mathbf{imK}}}_{{\mathbf{f}}}}$

where:

**i** = van’t Hoff factor

**m** = molality of the solution (in m or mol/kg)

**K _{f}** = freezing point depression constant (in ˚C/m)

Recall that the ** molality of a solution** is given by:

$\overline{){\mathbf{molality}}{\mathbf{=}}\frac{\mathbf{moles}\mathbf{}\mathbf{solute}}{\mathbf{kg}\mathbf{}\mathbf{solvent}}}$

**For this problem, we need to do the following:**

* Step 1:* Calculate for ΔT

* Step 2:* Determine the molality of the solution.

* Step 3:* Calculate the mass of C

Freezing Point Depression

Freezing Point Depression

Freezing Point Depression

Freezing Point Depression