Energy of Phase Changes

Objective: 

To determine the energy involved in phase changes of several substances.

Background:

    Consider heating a solid:  as the solid is warmed, energy (from some source) is "put into" the solid, and the solid gains energy.  As heating is continued, the warming solid reaches its melting point.  As more energy is added to the substance, the substance changes phase, but during this phase change, the temperature stays the same, resulting in a liquid at the melting point temperature.

     Of course, this liquid may be warmed as more energy is put into it, and if heating (energy input) is continued, the liquid will eventually reach its boiling point, at which another phase change--from liquid to gas--occurs.  This phase change requires energy, again with no change in temperature, but results in a more energetic phase, a gas.  It is possible, with the continued application of heat energy, to warm the gas.  Some substances, with application of heat energy to the solid, change from solid directly to gas: this is called sublimation.  

      It is possible to place two substances in contact with each other (perhaps by simply combining them in a well-insulated container) such that heat energy is transferred from the first (warmer) substance to the second (cooler) substance with minimal energy loss to the surroundings.  The energy change of the first substance can be calculated from its mass m, temperature change DT, and specific heat s (which is 4.184 J/g C, for water), according to the equation

q = msDT

J  =   g  x  [ J / ( g  x  C ) ]  x  C

If the system is indeed well-insulated, then the energy change of the second substance is assumed to be equal, but opposite in sign, to the energy change of first substance ("opposite in sign" refers to a + sign versus a - sign, or vice versa).  If the amount of the second substance is known, the calculation can be taken further: the energy involved with the phase change can be stated in units of energy/g or energy/mol (typically expressed as J/g or kJ/mol).  

    The energies involved with phase changes are:  

Heat of... Symbol Phase Change
Fusion DHfus Freezing or melting
Vaporization DHvap Boiling or condensation
Sublimation DHsub Solid to gas or
Gas to solid

 The magnitudes of these changes are characteristic for each substance under the conditions of the experiment.  

     Note the difference in units.  Specific heat has units of J/g C, and a temperature change is involved.  The heat involved in a phase change has units of J/g (or kJ/mole), and the temperature does not change!

    The measurement of heat involved in reactions is called calorimetry (the measurement of calories, another unit of heat energy).  It has many applications.  For example, the energy content of foods can be measured.  Listed below are a number of web sites describing this useful technique and how it is applied in some research projects.

http://itl.chem.ufl.edu/2045/lectures/lec_9.html

http://www.sciencebyjones.com/energy_content_of_food.htm

[You might find it interesting that for many elements, the product of the specific heat times the atomic weight has approximately the same value!  This was used to determine the approximate atomic weight for some elements!  (Petit and Dulong, Annales de Chimie et de Physique 10 pg 395, 1819  and Magie, Wm F., A Source Book in Physics McGraw-Hill NY 1935 pg 178)]

 

Pre-laboratory

(1)  During evaporation, molecules escape from the surface of a liquid and enter the gas phase.  This physical change requires an energy input: a molecule cannot become gaseous unless it has sufficient energy to break free of its interactions with neighboring molecules.  Is evaporation an exothermic or endothermic process?  (Hint: Section 6.2 of your Chang Chemistry textbook describes exothermic versus endothermic processes.)

 

(2)  Hold a small piece of ordinary ice in your hand and allow it to melt completely.  Record your observations.  Based on this experiment, do you think the melting process exothermic or endothermic?  Explain.

 

(3)  A 46.6-g sample of warm water was combined with a sample of liquid nitrogen in a Styrofoam cup calorimeter.  The liquid nitrogen sample rapidly evaporated, leaving only cool water in the calorimeter.  Careful temperature measurements showed that the water cooled from an initial temperature of 74.60C to a final temperature of 8.50C.  Calculate the amount of heat released by the water.  Express your answer in units of kilojoules.  (Hint: In addition to the Background information above, consult section 6.5 in your Chang Chemistry textbook if you need help.)

 

(4)  As mentioned in the Background section above, DHvap refers to the enthalpy or "heat" change that accompanies the vaporization of 1 mole of a liquid substance.  In a Styrofoam cup calorimeter experiment, a Chem 184 student found that 10.9 kJ of energy  was consumed by the vaporization of a 56.0 g sample of N2 (liquid nitrogen).  Use the student's experimental data to calculate a value for DHvap for N2.


Safety 

    At atmospheric pressure liquid N2 boils at -196C and dry ice (solid CO2) sublimes at -78C.  These substances are very cold and can quickly cause frostbite to exposed skin, so care must be taken in handling these substances.

   Take care to keep the thermometer and Vernier temperature probe from direct contact with the liquid nitrogen and the dry ice.

    Of course, wearing of goggles is imperative at all times in the laboratory.

 

Instructions for Part 1 

Reagents:  Ice

You will use 2 pairs of nested styrofoam cups.  Mark and obtain the mass of each nested pair, so that you will be able to determine the mass of water, etc. that you add to them. Prepare a data table in you lab notebook.

            Heat approximately 250 mL of water to 55 C - 65 C.  When the water has been heated to the desired temperature, pour approximately 60 mL into one of the nested pairs and obtain and record its mass.  Preheating or precooling the nested cups with a small amount of the water is recommended to avoid heat loss from the water to the container. 

            With your second cup pair on the balance, add approximately 20 grams of ice and record the mass. (Take your glass thermometer to the balance with you.) After first confirming the actual temperature of your water, pour the ice into the nested cup pair holding the water. 

Return to your work station and immediately initiate data collection using the Vernier temperature probe. (For instructions on use see the procedures page.) Stir the contents gently with the probe.  When all the ice has melted and the temperature has reached a minimum and begins to rise, you may stop the data collection.  By clicking on the STAT button, you can obtain the minimum temperature. Each group must collect at least three data sets for the heat of fusion.  (Remember in your calculations that the water formed from the melted ice is being warmed!) Find an average value for DHfus of ice and express it in units of J/g (i.e., the heat of fusion), and also in units of kJ/mol, the molar heat of fusion.  Add your results to the class data on the board.  Record the data in your lab notebook.

 

Instructions for Part 2

Reagents:  Liquid nitrogen

Use the same pairs of cups as in Part 1 and be certain that they are thoroughly dry.  Heat approximately 250 mL of water to 55 C - 65 C as before.  When the water has been heated to the desired temperature, pour approximately 60 mL into one of the nested pairs. Obtain and record its mass as you did in the first Part of the lab.  Again, preheating or precooling the nested cups with a small amount of the water or liquid nitrogen is recommended to avoid heat loss from the water to the container, and also to reduce premature loss of liquid nitrogen to the atmosphere. 

            With your second cup pair on the balance, add approximately 40 grams of liquid nitrogen and record the mass. (Take your glass thermometer to the balance with you). After first confirming the actual temperature of your water, pour the nitrogen into the nested cup pair holding the water.  Liquid nitrogen changes phase quickly. In order to prevent significant loss of mass, make the transfer to the water as soon as you remove the cups containing the liquid nitrogen from the balance - dont wait until you get back to your bench.  A cloud of extremely cold vapor will form above the container.  Fan this away to avoid cooling the contents of the cup.  

            Return to your computer and immediately initiate data collection using the Vernier temperature probe as you did in Part 1.  Stir the contents gently with the probe.  When all the liquid nitrogen has evaporated and the temperature has reached a minimum and begins to rise, you may stop the data collection.  By clicking on the STAT button, you can obtain the minimum temperature. Each team must collect at least three data sets for liquid nitrogen.  Find average values for DHvap of liquid nitrogen and express in units of J/g (i.e., the heat of vaporization), and in kJ/mole, the molar heat of vaporization.  Add your results to the class data on the board.  Record the data in your lab notebook.

 

Instructions for Part 3

Reagents:  Dry ice (solid carbon dioxide)

Repeat the procedure in Part 2, using 15 grams of crushed dry ice.  (Be sure that all the dry ice is transferred!)  Take care to ensure all the dry ice sublimes, and fan away the cold vapors from above the liquid surface.

Perform the experiment three times, and find the average value DHsub for dry ice in J/g and kJ/mol.

Add your results to the class data on the board.  Record the data in your lab notebook.

 

Procedures

    A measured amount of warm water will be placed into a pair of nested Styrofoam cups, and the temperature measured with a glass thermometer.  A weighed amount of a cold substance  will be added to the water and the temperature of the rapidly cooling water will be monitored and recorded using a Vernier temperature probe.  Using the temperature change of the water, its mass, and its heat capacity (4.184 J/g C), the heat lost by the water--and thus, the heat gained by the second substance--can be determined.  This energy change, together with the mass of the second substance, can then be used to determine the heat associated with the phase change in J/g and kJ/mol.  

    The temperature range used today will be approximately 20 C 70 C. Temperature will be measured initially with a glass thermometer and then monitored with the Vernier temperature probe.  Click here to obtain the procedure for setting up and using your computer and for data acquisition. You may wish to open Exp 18 from the Chemistry with Computers Experiment files of  Logger Pro.  

 

Questions to Consider

Of the three processes you investigated today, which involved the greatest amount of heat per mole for the phase changes involved?

Are the phase changes investigated exothermic or endothermic?

Which of the materials studied floated in water?

Does the amount of heat involved depend upon the amount of material being melted, vaporized, or sublimed?

Do you think the amount of heat involved depends upon the substance being vaporized/sublimed?

Which phase change requires the most energy?

What intermolecular changes are occurring?

What average values for DHfus , DHsub and DHvap were obtained by your group? How do these compare to the class results?

Discuss possible sources of error in this experiment.  How could you eliminate or minimize these sources of error?

This experiment was developed by Alfred Lata, Clark Bricker, and Susan Mason, with modifications by Roderick Black.  (See Burgstahler, A.W.; Hamlet, P. The Physics Teacher 1990, 28, 544-5 and Burgstahler, A.W.; Bricker, C.E., Journal of Chemical Education 1991, 68, 332-3.)