Freezing point depression pre lab answers

FREEZING POINT OF SOLUTIONS

I. PURPOSE In this experiment, you will determine the freezing point of water and aqueous solutions of methanol (CH

3 0H), a

non-electrolyte, and odium chloride (NaCl), an electro- lyte. With this information the molality of each solution will be calculated.

II. BACKGROUND All pure substances, elements, and compounds possess unique physical and chemical properties. Just as one hu- man being can be distinguished from all other by certain characteristics-such as fingerprints and DNA-it is also possible to distinguish any given compound from among the millions that are known.

Melting point and the boiling point are easily determined physical properties that are very useful in identifying a substance. Consequently, these properties are almost always recorded when a compound is described in the chemical literature (textbook, handbooks, journal articles, etc.). The freezing and melting of a pure substance occurs at the same temperature, measured when the liquid and solid phases of the substance are in equilibrium. When energy is being removed from a liquid in equilibrium with its solid , the process is called freezing; when energy is be- ing added to a solid in equilibrium with its liquid, the process is called melting.

Freezingenergy(-) Liquid ====:============::=: Solid Mel ring energy(+)

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86 EXPERIMENT 10

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A. MELTING AND FREEZING POINTS OF PURE SUBSTANCES When heat is removed from a liquid, the liquid particles lose kinetic energy and move

more slowly causing the temperature of the liquid to decrease. Finally, enough heat is removed and the particles move so slowly that the liquid becomes a solid; often a crys- talline solid. The temperature of the freezing point is different for different substances.

The amount of energy removed from a liquid to freeze, is equal to the amount of

energy added to its solid co melt. Thus, depending on the direction of energy flow, this equilibrium in temperatures is called the melting point or the freezing point.

B. FREEZING POINT OF IMPURE SUBSTANCES When a substance (solvent) is uniformly mixed with a small amount of another sub- stance (solute), the freezing point of the resulting solution (an "impure substance") will be lower than that of the pure solvent. For example, the accepted freezing point

for pure water is 0.0°C. Adding salt to water may make it freeze at temperatures as low as -21 °C depending on the amount of salt added to rhe water. This is the same concept when ethylene glycol, antifreeze, is added co the water in a car radiator to lower the freezing point of the water.

Melting point and freezing point data are of great value in determining the identity and/or purity of substances, especially in rhe field of organic chemistry. If a sample of a compound melts or freezes considerably below the known melting point of the pure substances, we know char the sample contains impurities, which have lowered

rhe melting point. If the melting point of an unknown compound agrees with that of a known compound, mixing the unknown compound with the known and determin- ing the melting point of the mixture can often confirm its identity. If the melting point of this mixture is the same as chat of the known compound, the compounds are identical. On the other hand, a lower melting point for the mixture indicates that the two compounds are nor identical.

A solution freezes at a lower temperature than does the pure solvent. This phenom- enon is called freezing point depression. The freezing point depression of a solution is a colligative property of the solution, which depends upon the number of dissolved particles in the solution. The higher the solute concentration, the greater the freezing point depression of the solution is. As shown in the following illustration:

FREEZING POINT OF SOLUTIONS

Freezing point pure liquid

Freezing point of solution

Pure liquid (blue)

Solution (red)

Time

Figure 10-1 . Freezing point plot of a pure solvent and a solution

The freezing point of the pure solvent is at a constant temperature but the freezing point of a solution slowly decreases. The decrease in the temperature of the freezing point is caused by the increase in solute concentration as the solvent freezes. The dis-

solved solutes can be non-electrolytes or electrolytes. Non-electrolytes are molecules that remain inract when they dissolve in water. Electrolytes are solutes that dissociate into ions, forming an electrically conducting solution. Non-electrolytes are molecules that remain intact when they dissolve in water and do not conduct electricity. The

equation describing the change in the freezing point depression (L1T1 ) from pure solvent to solution is:

LiT=T - T f f (puro olvcn1) f (solu1ion)

/J,. 1j > O; subtract the lower solution T_rfrom the higher solvent 1j

To find /J,. T_r, you need to consider if the solution you are working with is an electro- lyte or non-electrolyte. If the solution is an electrolyte then you need to use the van 't Hoff factor in your equation:

Where:

Li T_r = fy m (non-electrolytes)

LiT1 = i K1m (electrolytes)

K1 = the molal freezing point depression constant of the solvent (°C/m); for water it's l .86°C/m

m = molaliry = moles of solute per kilogram of solvent

= the number of dissolved particles (van 't Hoff factor)

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88 EXPERIMENT 10

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Ill. SAFETY PRECAUTIONS • Safety glasses are always required as long as anyone in the lab is still performing

laboratory work!

• An apron or lab coat and closed-toe/closed-heel shoes are required

• No food or drinks in the lab at any time.

• Methanol is flammable and is a poison if taken internally. If you get it on your skin, rinse it off with water as soon as possible. Methanol will dissolve nail polish, artificial fingernails, and can ruin some plastics.

• A broken thermometer is considered hazardous wasre. If the thermomerer is filled with mercury, it's highly toxic. If you break a thermometer, inform your TA right away.

• Wash your hands before you leave.

IV. MATERIALS AND REAGENTS

2

EQUIPMENT (communal equipment)

Thermometers

QT

20g

0.50 g

0.50 g

REAGENTS

Rock salt

Methanol

Granular salt

EQUIPMENT (from your

locker)

10 ml Graduated cylinder

250 ml Beaker

6" Test tube

~10 ml/ group WASTE

V. EXPERIME NTAL PROCEDURE You will be working with a lab partner on this experiment.

A. FREEZING POINT DETERMINATION OF SO LUT ION S

DISTILLED WATER (SO LVENT) 1. Obtain a 6" rest rube (be sure that it is clean) , a 250 mL beaker, and from the

communal equipment, two thermometers.

2. Using the 10 mL graduated cylinder, measure 8 ml of deionized warer and place ir into the 6" rest rube. Position one of the rhermomerers in the resr rube so rhar rhe end of rhe thermometer (bulb) is about 1.5 cm from the borrom of the rest rube (see Figure 10-2). Position the rhermomerer so char the scale can be read e:tSily. Make sure chat the thermometer's bulb is immersed near rhe center of the solution inside the rube. Place rhe rest rube in an available beaker. Remember thermometers can roll off the bench tops and break. To prevent this, make sure your thermometer is in a secure location to avoid rolling off of the bench top.

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FRE EZING POI NT OF SOLUTIONS 89

3. Fill a 250 ml beaker about halfway full of crushed ice and add about 20 g of

rock sale and mix it. Place the ocher thermometer in chis beaker and check the temperature; it should be closer to - l0°C.

4 . C reate a data cable and open the stopwatch application on your phone or check

the second hand of the clock on the wall. One person in your group will be read-

ing the temperature and the ocher will collect data every 20 seconds.

5. Before introducing the test tube to the ice bath, record the temperature of the distilled water every 20 econds in your data cable until it has reached thermal equilibrium (five consecutive points char are +/- 0.2°C). Lower the rest cube into the ice bath so that all of the distilled water in the rube is below the surface

of the ice water. Start collecting the temperature every 20 seconds. Make sure co

position the thermometer in a way that it is easy to read. Hold the beaker with one hand and the rest cube with the other and swirl the test cube continuously co have a homogeneous temperature.

6. Once the temperature reaches around 1 °C, lift che thermometer and stir the solu- tion in the test rube. You will notice a chin layer of ice form on the edge of the

test cube. After some amount of the ice or slush is formed inside of the test rube, remove it from the ice bath and stir vigorously. Make sure you note the constant temperature during the time that ice and liquid water are both present. This should be zero; cake three more readings and wait until the last crystal melts. This is the freezing point of pure water.

7. After completing the temperature readings, continue with a second trial, just letting the water reach room temperature. Ir will help to place the rest cube in a beaker containing water at room cemperacure.

Ice, water, and rock salt

Figure 10-2.

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90 EXPERIMENT 1 O

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WATER WITH METHANOL (NON -ELECTROLYTE) 8. ext, let the water stabilize to room temperature and then add 0.500 g of metha-

nol (CH 3 OH). Under the hood, there is a bottle dispenser. Use this as demon-

strated by your TA; handle this very carefully. This is calibrated co dispense 0.25 g

of methanol; the density of methanol is 0.792 g/mL.

9. Repeat steps 3-7; use the same ice in the beaker, and if necessary, remove some of the water and add more sale (add about IO g) and ice (co fill half of the beaker) until you have a temperature closer co -10°C in the beaker. Record the tempera- ture of your rest tube in your lab notebook.

10. Clean the test rube.

WATER WITH SODIUM CHLORIDE (ELECTROLYTE) 11 . Measure another 8 mL of water and this time, add approximately 0.500 g of

granular salt. Note: Do not add rock salt to the test tube.

12. Repeat steps 3-7; use the same ice in the beaker, and if necessary, remove some of the water and add more salt (add about 10 g) and ice (co fill half of the beaker) until you have a temperature closer to - 10°C in the beaker. Record the tempera- ture of your test tube in your lab notebook.

VI. GRAPHING TEMPERATURE DATA Graph the three sets of data in your lab report and identify the freezing point depres- sion of each solution by identifying on the graph the freezing point of solution and the freezing point of solvent.

VII. CALCULATIONS 1. Calculate the theoretical molality of the methanol and sodium chloride solu-

tions.

2. Calculate the experimental molality of the methanol and sodium chloride solu- tions.

3. Calculate the percent difference between the experimental and theoretical molal- icy of the methanol and the sodium chloride solutions.

VIII. DISCUSSION/QU ESTIONS 1. Discussion: Discuss the accuracy, precision, systematic errors, and random errors

of your data.

2. When the solid and liquid phases are in equilibrium, which phase contains the greater amount of energy? Explain.

3. Is acetic acid an electrolyte? \Vhy or why not?

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FREEZING POINT OF SOLUTIONS

REFERENCES "Freezing Point of Acid Solutions" Sci Am Scienrific American 16.397supp (1983):

6333. Web. From http:/ /www.lahc.edu/classes/chemisrry/arias/Exp%2012%20

- 0 o20Freezing%20PoinrF1 l .pdf

"Colligarive Properties: Freezing Point." Chemical Thermodynamics at a Glance (n.d.): 168-73. Web. From http:/ /themalloryfamily.ner/chemisrry/cheml 021/

Exp%200%20-%20Colligative%20Properties.pdf

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