Freezing Point DepressionBY: Dr. Robert D. Craig, Ph.D.

Concept of phase diagram

Concept of phase diagram

For carbon dioxide

A supercritical fluid

• A supercritical fluid is any substance at a temperature and pressure above its critical point, where distinct liquid and gas phases do not exist. It can effuse through solids like a gas, and dissolve materials like a liquid. In addition, close to the critical point, small changes in pressure or temperature result in large changes in density, allowing many properties of a supercritical fluid to be "fine-tuned

To extract caffiene-

Freezing Point Depression

• Colligative properties include: relative lowering of vapor pressure; elevation of boiling point; depression of freezing point and osmotic pressure. Measurements of these properties for a dilute aqueous solution of a non-ionized solute such as urea or glucose can lead to accurate determinations of relative molecular masses

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Depression of the Freezing Point of a Solvent

• We need two pieces of information to calculate the depression of the freezing point of the solvent in a solution containing a nonvolatile nonelectrolyte:

• The molal concentration, m, of the solute in the solution (Sometimes it may be necessary to calculate this concentration from other information.)

• The freezing point depression constant, Kf, for the solvent.

• We use the following equation to calculate a freezing point depression. (Note: This equation assumes the freezing point depression constant has a positive value.)

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• As is demonstrated by the phase diagram above, adding a solute to a solvent lowers the freezing point and raises the boiling point; it also lowers the vapor pressure.

• The new freezing point of a solution can be determined using the colligative property law:

• ΔTf = kf m

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• The change The change in freezing point is equal to the molal freezing-point constant times the molality of the solution. The molal freezing-point constant used is the constant for the solvent, not the solute.

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• In this experiment, the molar mass of sulfur will be determined using the colligative property law. The freezing point of naphthalene will be determined

• experimentally; then a controlled solution of naphthalene and sulfur will be made, and the freezing point of that solution will be determined. The difference in freezing point can be used in the colligative property law to determine the experimental molality of the solution, leading to a calculation

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• The freezing temperature is difficult to ascertain by direct visual observation because of a phenomenon called supercooling and also because solidification of solutions usually occurs over a broad temperature range. Temperature-time graphs, called cooling curves, reveal freezing temperatures rather clearly.

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• The cooling curve will look like the

• one below in figure 19.2:n of molecular weight.

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• In order to minimize supercooling, the solution will be stirred while freezing.

• To determine the molar mass of a substance, one must simply divide the grams of substance by the number of moles of substance present. All of these values will be determined experimentally.

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• Procedure:• Part A – Cooling Curve for Pure Naphthalene• 1. A large test tube was weighed to the

nearest .01 g using a standard laboratory balance.

• Approximately 15 to 20 grams of naphthalene was added to the test tube. The test tube

• was weighed again using the standard balance.

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• 3. The 600-mL beaker was nearly filled with water. It was heated to about 85°C. The test

• tube was clamped in the water bath as shown in Figure 19.3 above. When most of the

• naphthalene had melted, the stopper containing the thermometer and stirrer was placed

• into the test tube. The thermometer was not allowed to touch the bottom or sides of the

• test tube

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• 4. When all of the naphthalene had melted, the test tube was removed from the beaker of boiling water. The test tube was placed into a wide-mouthed bottle with some paper towels at the bottom. The temperature reading from the thermometer was recorded every 30 seconds.

• The naphthalene was stirred using the wire stirrer to ensure even freezing.

• When the temperature remained constant for several readings, the naphthalene was

• allowed to cool without further temperature readings.

freezing point depression

• We use the following equation to calculate a freezing point depression. (Note: This equation assumes the freezing point depression constant has a positive value.)

• Freezing Pointsolution = Freezing Pointsolvent - ΔTf

freezing point depression

• Freezing Pointsolution = Freezing Pointsolvent - ΔTf

• where

• ΔTf = molality * Kf * i, (Kf = cryoscopic constant, which is 1.86°C kg/mol for the freezing point of water,; i = Van 't Hoff factor)

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the enthalpy change

the enthalpy change• If the enthalpy change of reaction is

assumed to be constant with temperature, the definite integral of this differential equation between temperatures T1 and T2 is given by his equation

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Freezing Point of a Solution

• The freezing point of the solvent in a solution containing a nonvolatile nonelectrolyte, Tsolution, may be found from the following information:

• ΔTF = KF · m · i

The freezing point of the pure solvent, • Tpure solvent. • The freezing point depression, T. • Note: The freezing point of a solution will always

be lower than the freezing point of the pure solvent.

Ideal solution-like ideal gas

• ΔTF =TF KF · m · I

• ΔTF, the freezing point depression, is defined as TF (pure

solvent) - TF (solution).

• KF, the cryoscopic constant, which is dependent on the properties of the solvent, not the solute. Note: When conducting experiments, a higher KF value makes it easier to observe larger drops in the freezing point. For water, KF = 1.853 K·kg/mol.[4]

• m is the molality (mol solute per kg of solvent)

van 't Hoff factor• i is the van 't Hoff factor (number of solute

particles per mol, e.g. i = 2 for NaCl).

Depression of Napthalene

Depression of Napthalene

 When a pure liquid cools its temperature decreases

Steadily until the freezing point is reached. Then its

Temperature remains constant as the liquid solidifies.

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At constant pressure, the pure liquid and pure solid can be in equilibrium at only one temperature , the freezing point of the compound. When all of the liquid has frozen, the temperature of the solid begins to decrease.

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