Chemistry Lab Report

CHEMISTY 135/150/175Chemistry Lab
Energy of Phase Changes[1]
Water at its triple point, where 3 phases coexist, looks unfamiliar—a bit like boiling water and ice. [2]
Coexistence of two phases, such as an ice cube in a glass of water, are much more common.
What energy is required to convert from one phase to another? How do we measure it?
DEPARTMENT OF CHEMISTRY
UNIVERSITY OF KANSAS
1
Introduction
Consider heating a solid: as the solid is warmed, energy from the source of heat is “put into” the solid, and
the solid gains energy. If the heating is continued, the warming solid will eventually reach its melting point and
convert to a liquid. To complete this transformation, more energy (e.g., in the form of heat) must be added to the
substance. However, as the substance changes phase, the added energy goes into accomplishing this
transformation, and the temperature stays the same. Thus, a single melting point temperature can be determined
at which the solid converts to the liquid. This is an example of a phase transition. The amount of heat required to
accomplish the solid-to-liquid phase change at the melting temperature is an enthalpy and is called the heat of
fusion.
Of course, if the resulting liquid is also heated its temperature will increase as more energy is put into it
until the liquid eventually reaches its boiling point, at which another phase change occurs, this time from a liquid
to a gas. Again, at this temperature, the continued heating provides energy for the liquid-to-gas phase transition
while the temperature remains the same. The heat, or enthalpy, required for the liquid-to-gas phase change is
known as the heat of vaporization. It is worth noting that the above assumes the substance has both solid-to-liquid
and liquid-to-gas transitions. Some substances convert directly from a solid to a gas when heated sufficiently, a
process called sublimation. The heat needed for this solid-to-gas transition is called the heat of sublimation.
In the processes described above, note that there are two steps. First, the substance is heated, generally by
placing it in contact with another substance that is warmer (i.e., at a higher temperature, T). Energy is then
transferred as heat from the warmer substance to the cooler substance (this continues until they reach the same
temperature). If the two substances are in a well-insulated container such that not much heat escapes to the
surroundings, the energy change of the initially cooler substance can be calculated from its mass m, temperature
change ΔT, and heat capacity or specific heat s, which is 4.184 J/(g oC) for water, according to the equation
(Eqn 1)
𝑞 = 𝑚𝑠∆𝑇
If the system is indeed well-insulated, then the energy change of the second, initially warmer substance is
assumed to be equal, but opposite in sign, to the energy change of first substance. Second, at the temperature of
the phase transition, the energy transferred as heat goes into accomplishing the change of phase while the
temperature remains constant. If it is the cooler substance undergoing the phase transition, the heat energy required
for the phase change comes from the warmer substance, which does change temperature as it loses heat. Thus, for
a substance undergoing a phase change at the transition temperature, the heat energy required can be obtained from
the change in temperature of its surroundings. 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 latter is the molar heat for the phase transition.
Note the difference in units. In this way, the energy required to change the phase of a substance can be measured
using a calorimeter by isolating the substance with, e.g., a warmer substance that provides the required heat. Then
measuring the drop in temperature of this warmer substance gives the heat required for the phase transition. The
energies involved with phase changes are:
Heat of …
Symbol
Phase Change
Fusion
ΔHfus
solid  liquid; melting
Vaporization
ΔHvap
liquid  gas; boiling
Sublimation
ΔHsub
solid  gas; sublimation
Note: The magnitudes of these phase changes are characteristic for each substance under the conditions of the
experiment.
2
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. Below is a link that further describes the application of this useful technique in some research
projects.
http://www.dairyscience.info/packaging-/119-labelling-determination-of-the-energy-content-of-food.html
Interestingly, 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 forsome
elements.[2]
Safety
Liquid N2 boils at -196°C (at atmospheric pressure). Solid CO2 (dry ice) sublimes at -78°C (at
atmospheric pressure). These substances are very cold and can quickly cause frostbite to exposed skin,
so care must be taken in handling them.
Procedure
In this three-part experiment, 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 undergoing
a phase transition 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
specific heat of 4.184 J/(goC), the heat lost by the water — and thus the heat gained by the cold substance — can be
determined. This energy change, together with the mass of the cold 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.
Part 1 – The Heat of Fusion of Water (Ice)
In this part of the experiment, you will use a simple calorimeter to determine the heat of fusion of ice.
1. Confirm that the Vernier LabPro interface box is connected to the computer via a USB cable. The stainlesssteel temperature probe should be connected to the CH1 port in the LabPro box.
2. Open LoggerPro software and configure the software for data collection by opening file number 18 from the
“Chemistry with Vernier” folder. If the real-time temperature reading is not displayed, ask for help.
3
3. Obtain 2 pairs of nested Styrofoam cups (i.e., 4 total cups). Label each nested pair, then measure and record
their masses separately in your notebook. In this way, you will be able to determine the mass of water and
other substances that added to them.
4. Prepare Table 1 in your lab notebook.
Table 1. Calorimetry Data for Melting of Ice
m[H2O(liquid)] (g)
m[H2O(solid)] (g)
Ti[H2O(l)] (°C)
Tf[H2O(l)] (°C)
ΔT[H2O(l)] (°C)
Trial 1
Trial 2
Trial 3
Note: The mass of the liquid water is calculated from the difference between the empty cup and the full cup. Be sure
to keep track of which pair of cups you use as not every pair weighs the same.
5. Heat approximately 250 mL of water to 55 °C – 65 °C.
It should be noted that the next steps (i.e., 6 – 12) should be done carefully and quickly to minimize the heat lost to
the surroundings. Also, have a glass thermometer at the ready to measure temperature.
6. Pour approximately 60 mL of the warm water into one of the nested pairs and record the mass.
7. Calculate the mass of liquid water by subtracting the pair of cups’ mass. Record the mass of the liquid water
in Table 1.
8. Remove the cups with liquid water from the balance and place the empty second cup pair on the balance. Tare
the balance.
9. Add approximately 20 grams of ice and record the mass in Table 1.
10.
Using the glass thermometer, record the initial temperature of the warm water, T i.
11.
Pour the ice into the nested cup pair holding the water.
12. Return to your workstation and immediately place the Vernier temperature probe into the calorimeter. Click
the “Collect” button at the top of the LoggerPro software window to initiate data collection. 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, Tf.
13.
Repeat steps 6-12 two more times for a total of three trials.
Part 2 – The Heat of Vaporization of Nitrogen
In this part of the experiment, you will use a simple calorimeter to determine the heat of vaporization of nitrogen.
1. Prepare a data table (label it table 2 and give it a descriptive title) in your lab notebook analogous to the one in
Part 1. The only difference will be instead of solid ice mass you have liquid nitrogen mass.
2. Again, heat approximately 250 mL of water to 55 ºC – 65 ºC.
3. 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 Part 1.
The next few steps require some coordination. Read the steps carefully before proceeding.
4. With your second cup pair tared on the balance and the warm water cups outside the balance, measure the initial
temperature of the water, Ti. Then, add approximately 40 grams of liquid nitrogen to the second cup pair and
4
record the mass. The mass of liquid nitrogen may slowly decrease. Be sure to record the mass just before the
cups are removed from the balance.
5. Then, carefully pour the nitrogen into the nested cup pair holding the water. Note: Liquid nitrogen changes
phase quickly. A cloud of extremely cold vapor will form above the container. Gently fan this away.
6. 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.
7. Repeat steps 2-6 two more times for a total of three trials.
Part 3 – The Heat of Sublimation of CO2 (Dry Ice)
In this part of the experiment, you will now use a simple calorimeter to determine the heat of sublimation of CO 2,
otherwise known as dry ice.
Repeat the procedure from part 2, except you this time you will use 15 grams of dry ice. As before, complete
process in triplicate (i.e., three trials). Note you will need to warm at least 200 mL of water to 55 ºC – 65 ºC before
beginning.
Calculations
In each part of the lab your group measured the heat transferred during a phase change. The following will
aid in converting that heat measured into the enthalpy, H, of the phase change.
Part 1 – The Heat of Fusion of Water (Ice)
1. Construct the following table in your notebook.
Table 4. Calorimetry Data for Melting of Ice
Mass of ice (g)
Moles of Ice
Trial 1
Trial 2
Trial 3
qmelt (J)
Hfus (J/g)
Hfus (kJ/mol)
Average
Standard Deviation
2. Place mass of ice from each trial into the table and then convert the mass of ice to moles of ice using
molar mass of water.
Calculating the heat of melting, qmelt, for ice can be difficult because not only does the ice melt, it also
then heats up to a final temperature (i.e., the ice melts and then the residual liquid warms up). This can be
expressed using the following equation:
𝑞
+ 𝑞
,
= −𝑞
(Eqn 2)
,
The heat of liquid cooling comes from the warm water in the calorimeter cooling down and the heat
of ice warming comes from the liquid ice warming up after melting. The equation can be expanded to
describe these in greater detail.
𝑞
+ (𝑚
∗ 𝑠 ∗ 𝑇 − 𝑇 ) = −(𝑚
∗𝑠∗ 𝑇 −𝑇 )
(Eqn 3)
3. Now, you have all the variables needed to calculate the heat of melting, qmelt. Work with your group
members and the TA to calculate heat of melting for trial 1 – 3 and fill in the table correspondingly.
4. Using the mass of ice and the heat of melting, calculate enthalpy of fusion (melting) in Joules per gram.
5
5. Then, use the moles of ice and the heat of melting to calculate the enthalpy of fusion (melting) in
kilojoules per mole. Note the units are now in kilojoules instead of joules.
6. Using excel, determine an average and standard deviation for Hfus.
Part 2 – The Heat of Vaporization of Nitrogen
1. Construct the following table in your notebook.
Table 5. Calorimetry Data for Vaporization of Dinitrogen
Mass of N2 (g)
Moles of N2
Trial 1
Trial 2
Trial 3
Hvap (J/g)
qvap (J)
Hvap (kJ/mol)
Average
Standard Deviation
2. Place mass of liquid nitrogen from each trial into the table and then convert the mass to moles using
molar mass of dinitrogen.
Calculating the heat of vaporization, qvap, is a little bit easier than ice melting because unlike ice,
liquid nitrogen vaporizes and leaves the calorimeter completely! This can be expressed using the following:
𝑞
(Eqn 4)
= −𝑞
Expanding the equation gives the following:
𝑞
= −(𝑚
(Eqn 3)
∗𝑠∗ 𝑇 −𝑇 )
3. Now, you have all the variables needed to calculate the heat of vaporization, qvap. Work with your
group members and the TA to calculate heat of vaporization for trial 1 – 3 and fill in the table
correspondingly. Ask questions where needed.
4. Using the mass of liquid nitrogen and the heat of vaporization, calculate enthalpy of vaporization in
Joules per gram.
5. Then, use the moles of dinitrogen and the heat of vaporization to calculate the enthalpy of vaporization
in kilojoules per mole. Note the units are now in kilojoules instead of joules.
6. Using excel, determine an average and standard deviation for Hvap.
Part 3 – The Heat of Sublimation of CO2 (Dry Ice)
1. Construct the following table in your notebook.
Table 6. Calorimetry Data for Sublimation of CO2
Mass of CO2 (g)
Moles of CO2
qsub (J)
Hsub (J/g)
Trial 1
Trial 2
Trial 3
Average
Standard Deviation
6
Hsub (kJ/mol)
2. Complete the table above for the sublimation of CO2 using data from your calorimetry data. Note that dry
ice sublimating is similar to liquid nitrogen vaporizing in that the carbon dioxide leaves the calorimeter
during the process. Thus, the calculation process for enthalpy of sublimation will be similar to that of the
enthalpy of vaporization of dinitrogen calculation.
Part 1-3 Literature Comparison
1. The literature values (i.e., the theoretical values) for the measured enthalpies are as follows: water’s Hfus
= 6.02 kJ/mol, dinitrogen’s Hvap = 5.56 kJ/mol, and carbon dioxide’s Hsub = 25.2 kJ/mol. Using the
average experimental values, calculate the % error in your measurements.
|𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑣𝑎𝑙𝑢𝑒 − 𝑒𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙 𝑣𝑎𝑙𝑢𝑒|
× 100 = % 𝑒𝑟𝑟𝑜𝑟
𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑣𝑎𝑙𝑢𝑒
Concepts to Discuss with Group Members
1. Is the phase change exothermic or endothermic? Do you think the amount of heat involved in a phase
change depends upon the identity of the substance being melted, vaporized, or sublimed?
2. Of the three processes you investigated today, which involved the greatest amount of heat per mole for the
phase changes involved?
3. What intermolecular changes are occurring during each of the three types of phases changes your group
investigated?
4. What are possible sources of error in this experiment? How could you eliminate or minimize each of these
sources of error?
Reference(s)
[1] Significant portions of this experiment were developed from the laboratory manual Experiments to Establish Foundations
of Chemistry I, by Alfred J. Lata and Clark E. Bricker, University of Kansas, 1995. (See also Burgstahler, A.W. & Hamlet,
P. The Physics Teacher 1990, 28, 544-5; Burgstahler, A.W. & Bricker, C.E., Journal of Chemical Education 1991, 68, 332-3;
and Spaeth, A. D. & Black, R.S., Journal of Chemical Education 2012, 89, 1078-9.)
[2] http://fphoto.photoshelter.com/image/I0000FcqGAZrJYLU accessed June 16, 2016.
[3] Petit and Dulong, Annales de Chimie et de Physique 10 pp. 395, 1819; and W. F. Magie, A Source Book in Physics,
McGraw-Hill: NY, 1935, p. 178
7
Glossary
Calorimeter
Heat of vaporization (heat of condensation); ∆𝐻vap
the experimental procedure used to measure the heat
released from or absorbed by a chemical reaction or
phase transition
the amount of heat required to convert one mole of a
liquid to a gas (or gas to liquid); because
vaporization and condensation are the reverse
process of each other, ∆𝐻vap = −∆𝐻cond; ∆𝐻vap is
usually reported
Coffee-cup calorimeter
a simple piece of equipment designed to measure ∆𝐻
for reactions or phase transitions occurring under
conditions that maintain a constant pressure
Heat of fusion (enthalpy of freezing); ∆𝐻fus
the amount of heat required to convert one mole of a
solid to a liquid (liquid to solid); generally,melting—
also called fusion—and freezing are the reverse
process of each other, and ∆𝐻fus =
−∆𝐻freeze; ∆𝐻fus is usually reported
Condensed phase
a phase of matter in which molecules are in constant
contact with neighboring molecules
Liquid
Fluid
a phase of matter characterized by individual
molecules being in constant contact with
neighboring molecules and the tendency of this
collection of molecules to assume the shape of the
vessel that contains them; the fluid condensed phase;
a phase of matter typically characterized by high
temperature and high pressure, or low temperature
and low pressure, relative to other possible phases
a phase of matter in which associations between
molecules do not constrain the overall shape of the
collection of molecules; any phase in which the
collection of molecules assumes the shape of the
vessel that contains them
Gas (vapor)
Phase change (phase transition)
a phase of matter characterized by large separations
between molecules and the tendency of the
collection of molecules to assume the shape of the
containing vessel; the fluid, non-condensed phase; a
phase of matter typically characterized by high
temperature and low pressure, relative to other
possible phases
the conversion of one phase to a different phase
Phase (of matter)
a region of a material, containing one or more kinds
of molecules, exhibiting the same physical properties
in a uniform way across the whole region; a region
of a material that is chemically uniform, physically
distint, and mechanically separable; “phase” is often
incorrectly conflated with “state of matter”; phases
of matter can include various kinds of mixtures and
can include the same state of matter with different
physical properties or forms (e.g., two different
kinds of solid phosphorus)
Heat of sublimation (heat of deposition); ∆𝐻sub
the amount of heat required to convert one mole of
solid to gas directly (gas to solid directly); because
sublimation and deposition are the reverse process of
each other, ∆𝐻sub = −∆𝐻dep; ∆𝐻sub is usually
reported.
8
Solid
a phase of matter characterized by individual
molecules being in constant contact with
neighboring molecules and the tendency of this
collection of molecules to adopt a fixed shape; the
non-fluid condensed phase; a phase of matter
typically characterized by low temperature and high
pressure, relative to other possible phases
9