CHEM&152
Spr ing 2009
INTERMOLECULAR FORCES and SOLID-LIQUID TRANSITIONS
Fill-in
Name _______________________
Pre-lab attached (p 10-11)
Stamp Here
Lecture Instructor __________________
Partner ______________________
Date _________________
Parts 1 and 2 of this experiment may be done on separate days.
Helpful info: Today you will be using a program called LoggerPro. It is a powerful program where
data can be acquired using probes, and sensors. The data can then be manipulated. Below is a key
for the shortcut menu displayed by LoggerPro. Some of these buttons will come in handy in this
lab and in the future!
http://www.phas.ubc.ca/~year1lab/p100/LoggerPro/LoggerPictures/Toolbar.jpg
Part 1: Evaporation and Intermolecular Attractions
Purpose Investigate the relationship of dispersion forces and hydrogen bonding forces in
intermolecular attractions.
Discussion
In this experiment, temperature probes covered with filter paper are placed in various liquids.
Evaporation occurs when the probe is removed from the liquid's container (see figure 1). This
evaporation is an endothermic process that results in a temperature decrease. The magnitude of a
temperature decrease is, like viscosity and boiling temperature, related to the strength of
intermolecular forces of attraction. In this experiment, you will study temperature changes caused by
the evaporation of several liquids and relate the temperature changes to the strength of
intermolecular forces of attraction. The temperature change is greater if more evaporation occurs
(weaker intermolecular forces). You will use the results to predict, and then measure, the
temperature change for several other liquids.
You will encounter two types of organic compounds in this experiment — alkanes and alcohols. The
two alkanes are pentane, C5H12, and hexane, C6H14. In addition to carbon and hydrogen atoms,
alcohols also contain the -OH functional group. Methanol, CH3OH, and ethanol, C2H5OH, are two of
the alcohols that we will use in this experiment. You will examine the molecular structure of alkanes
and alcohols for the presence and relative strength of two intermolecular forces—hydrogen bonding
and dispersion forces. Dispersion forces exist between any two molecules and generally increase as
the molecular weight of the molecule increases. A hydrogen bond can occur when, in one molecule,
a hydrogen atom is bonded directly to an N, O, or F atom (the donor) and that hydrogen is
attracted to a lone pair on an N, O, or F atom in another molecule (the acceptor). Both a donor and
an acceptor must be present for a hydrogen bond to occur.
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FIGURE 1
Caution: In all experiments today keep liquids away from the computers and keyboards. Be
very careful when stirring with direct-connect temperature probes. Clean them gently with
water and wipe gently/blot with paper towels. Each temperature probe costs approx. $25.
Materials **Please keep the organic solutions in the bottles**
Computer
Serial Box Interface or ULI
Logger Pro
two Vernier Temperature Probes
6 pieces of filter paper (2.5 cm X 2.5 cm)
2 small rubber bands
masking tape
methanol, CH3OH
ethanol, C2H5OH
1-propanol, C3H7OH
1-butanol, C4H9OH
pentane, C5H12
hexane, C6H14
sodium thiosulfate pentahydrate
Procedure
1. Obtain and wear goggles. Collect the necessary reagents and materials.
2. Prepare the computer for data collection by opening the Vernier Software and LoggerPro. You
will need to then find the appropriate experiment using the pathway:
File  Open 
Experiments  Chemistry with Computers  09 Evaporation. You may need to confirm the
identity of the temperature probes that you are using in the Sensor Confirmation dialog box if it
pops up when you open the program. You should see a temperature scale on the y-axis and time
on the x-axis.
3. You will need to set the time scale for data collection. Under the Experiment menu, select Data
Collection. Change the Length to 400 seconds and click “OK”. Then rescale the horizontal axis
to 0 to 400 seconds by clicking once on the lowest number and changing it to 0, and clicking on
the highest number and changing it to 400. The vertical axis of the graph will have temperature
scaled from 5 to 30°C, change this to read 0 to 30° by clicking once on the minimum and
changing it to 0.
4. Wrap Probe 1 and Probe 2 with square pieces of filter paper secured by small rubber bands as
shown in Figure 1. Roll the filter paper around the probe tip in the shape of a cylinder. Hint: First
slip the rubber band up on the probe, wrap the paper around the probe, and then finally slip the
rubber band over the wrapped paper. The paper should be even with the probe end.
5. Stand Probe 1 in the ethanol container and Probe 2 in the 1-propanol container. Make sure the
containers do not tip over. The liquids are in bottles that will be reused by different groups so
be sure not to contaminate them.
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6. Prepare 2 pieces of masking tape, each about 10-cm long, to be used to tape the probes in
position during step 7.
7. After the probes have been in the liquids for at least 30 seconds, begin data collection by clicking
Collect . Monitor the temperature for 15 seconds to establish the initial temperature of each liquid.
Without hitting stop, simultaneously remove the probes from the liquids and tape them so the
probe tips extend 5 cm over the edge of the table top as shown in Figure 1. Make sure that the
filter paper goes to the tip of the probe, but not past the tip.
8. When both temperatures have reached minimums and have begun to increase, click Stop to
end data collection. Click the Statistics button, , then click OK to display a box for both
probes. Record the maximum (T1) and minimum (T2) values for Temperature 1 (ethanol) and
Temperature 2 (1–propanol) in the table.
9. For each liquid, subtract the minimum temperature from the maximum temperature to determine
∆T, the temperature change during evaporation.
10. Roll the rubber band up the probe shaft and dispose of the filter paper in a beaker your instructor
has placed in the hood.
11. Based on the ∆T values you obtained for these two substances, plus information in the Pre -Lab
exercise, predict the relative size of the∆T value for 1 -butanol. Compare its hydrogen-bonding
capability and molecular weight to those of ethanol and 1-propanol. Record your predicted∆T,
then explain how you arrived at this answer in the space provided. Do the same for pentane. It is
not important that you predict the exact∆T value; simply estimate a logical value that is higher,
lower, or between the previous ∆T values. (e.g. > 4°C or < 8°C)
12. Test your predictions in Step 10 by repeating Steps 3-9 using 1-butanol for Probe 1 and pentane
for Probe 2.
13. Based on the ∆T values you have obtained for all four substances, plus information in the Pre Lab exercise, predict the ∆T values for methanol and hexane. Compare the hydrogen -bonding
capability and molecular weight of methanol and hexane to those of the previous four liquids.
Record your predicted ∆T, then explain how you arrived at this answer in the space provided.
14. Test your prediction in Step 12 by repeating Steps 3-9, using methanol with Probe 1 and hexane
with Probe 2.
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Data Table
Substance
Tmax
(°C)
Tmin
(°C)
∆T (°C)
(Tmax–Tmin)
ethanol
Predicted
∆T (°C)
Explanation
1-propanol
1-butanol
pentane
methanol
hexane
Questions
1. Take a moment to look at your ∆T values. In the space below, organize your samples from
smallest ∆T to largest ∆T for your alkanes, and again for your alcohols:
Alkanes: ______________________ < ______________________
Alcohols: ________________< ________________< ________________ < ________________
2. Discover what the ∆T “tells” you about how readily a sample evaporates. Circle the appropriate
response to each statement below.
Large ∆T values mean that the sample
does
does not
evaporate readily.
Small ∆T values mean that the sample
does
does not
evaporate readily.
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3. Correlate your answers to how readily a sample evaporates to the strength of the intermolecular
forces in that molecule. (again, circle the appropriate response to complete the statement.)
If the sample evaporates readily, then the relative strength of the IMFs in the molecule is: (circle
one)
strong
weak.
If the sample does not evaporate, then the relative strength of the IMFs in the molecule is: (circle
one)
strong
weak.
4. Looking at the formulas for your alkanes, what is the only difference between formulas?
5. What IMF(s) are present in the alkanes (list all present)?
6. Looking at the formulas for your alcohols, what is the only difference between the formulas?
7. What IMF(s) are present in the alcohols (list all present)?
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Using your answers to the guided questions on page 5, answer the following questions:
1. Using Excel, plot a graph of∆T values of the four alcohols versus their respective molecular
weights. Plot molecular weight on the x-axis and ∆T on the y-axis. Label each point with the
identity of the alcohol. From the graph, comments on the relationship between the strengths of
intermolecular forces and a molecule’s molecular weight:
As molecular weight increases, the strength of a molecule’s IMF ______________________.
Explanation:
2.
Two of the liquids, pentane and 1-butanol, had nearly the same molecular weights, but
significantly different ∆T values. Based on their intermolecular forces, explain why there was a
difference in ∆T values of these substances. Be specific.
3.
Which of the alkanes studied has the stronger intermolecular forces of attraction? The weaker
intermolecular forces? Explain using the results of this experiment. Include in your
explanation comments about the strengths of their IMFs.
4.
Which of the alcohols studied has the strongest intermolecular forces of attraction? The weakest
intermolecular forces? Explain using the results of this experiment. Include in your explanation
comments about the strengths of their IMFs.
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Part 2: A Study of Solid-Liquid Transitions
Purpose Study the relationship of temperature to the transition of matter between the liquid and
solid states.
Discussion
Every compound has certain physical and chemical properties that are unique to that compound.
Two physical properties that are commonly recorded for a compound are boiling point and freezing
or melting point. Today, you will study the latter.
The melting and freezing points for a compound occur at the same temperature. This temperature is
the point where the compound's solid and liquid states exist in equilibrium with each other with no
change in temperature (at constant pressure). If the compound is turning to a liquid we talk about a
melting point, and if it is turning solid we call it the freezing point.
freezing point
Liquid 
→ Solid
Liquid ←

 Solid
melting point
As long as both the solid and liquid states are existing together, the temperature remains constant.
However, as soon as one of the states disappears, the temperature will change. All solids absorb heat
as they melt.
The heat of fusion is the amount of heat required to melt a unit quantity of substance (usually 1 g or
1 mole) with no change in temperature. For water this value is 6.01 kJ/mole or 79 cal/g. We know
that at the melting point no change in temperature occurs. If the solid is absorbing heat and no rise
in temperature is occurring, how then is this heat energy being used? In order to understand this we
need to first discuss the crystal properties of solids.
Most solid substances are crystalline. This means that in the solid state the atoms, ions or molecules
of the compound are arranged in a definite three-dimensional order. They are held in this order by
intermolecular forces holding the molecules (or atoms/ions) together. When we begin to heat a
crystalline solid, the molecules (or atoms/ions) take up the heat energy until they finally gain enough
energy to overcome the attractive forces holding them in the crystal structure. When a molecule
gains sufficient energy it breaks free of the crystal and becomes part of the liquid state. If the
heating is continued, eventually all the molecules in the solid will gain enough energy to break free
and the solid turns to liquid. At this point, any additional heat energy will raise the temperature of
the liquid.
Conversely, if no additional heat is added to the liquid it will begin to crystallize. As it crystallizes,
the liquid gives off heat energy. This heat energy keeps the temperature constant (i.e. at the
freezing point) until all the liquid crystallizes. Then the temperature begins to drop.
Sometimes it is possible to cool a liquid using an ice bath to a temperature that is below its normal
freezing point. A liquid existing in this state is said to be supercooled. This state is not stable and
crystallization will occur if the solution is stirred or if a "seed" crystal is added. The seed crystal
induces crystallization by providing a pattern of the crystal structure for the liquid compound to line
up with. As soon as crystallization begins, heat energy is given off and the temperature rises to the
normal freezing point and remains there until all the liquid has solidified.
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Procedure
In the experiment today you will measure the freezing point of a liquid and the freezing point of a
supercooled liquid using the same temperature probes you used in Part 1. One partner should start at
the Experimental Set-Up while the other does the Computer Set-Up.
Computer Set-Up
1. Verify that the computer is still set up for Experiment 9.
2. Under the Experiment menu open Data Collection. Choose the Collection tab. Change Length
to 30 minutes and collect at a sampling speed of 5 per minute. Press “Done”.
3. Change the temperature scale on the y-axis to read 0 to 100°C by clicking on the minimum and
changing it to 0 and then the maximum and changing it to 100. You should change the x-axis to
read from 0 to 30 minutes.
Experimental Set-Up
1. Obtain two test tubes from the lab bench which each contain about 10 g of hydrated sodium
thiosulfate crystals (Na2S2O3•5 H2O).
2. Fill a 250 mL beaker about 1/2 full with cold (use 3-5 ice cubes) water for use later. On a small
piece of weighing paper, obtain two “seed” crystals of sodium thiosulfate pentahydrate.
3. Do this part at your lab bench, not next to the computers. Disconnect the temperature probes
from the interface and insert them into the test tubes. Heat each test tube in a hot water bath,
stirring gently with the temperature probes until the crystals in both tubes are melted. If the
sodium thiosulfate hydrate begins to crystallize, you must reheat the test tube.
Be careful to keep the wires from the probes away from the hot plate or the burner. You may
use a clamp on the ring stand to suspend the wires.
Be careful melting the samples, monitor them closely. No crystals should be present once
melted, however, if the crystals are heated too hot, dehydration occurs and the melting point
changes. Do not contaminate the materials with water or other substances.
The next few steps must be done smoothly and quickly in order, MAKE SURE YOU READ
AHEAD AND ARE PREPARED!
4. Clamp the test tubes to a ring stand. Move the ring stand with the test tubes clamped onto it, to
the computer and reconnect each temperature probe. The temperature of both should be about
60°C. If the liquid begins to crystallize you will need to reheat it gently. You may do this with
the temperature probe in the test tube by unplugging it from the interface box. Be careful not to
heat the wires.
5. Press
Collect
to begin to collect data.
6. After about 30 seconds, raise the beaker of cold tap water around one of the test tubes (#2). This
will be your supercooled sodium thiosulfate pentahydrate.
7. When the temperature of the supercooled liquid reaches about 30°C (do not stir the supercooling
tube or you may cause crystallization), remove the beaker of water first, then drop a “seed”
crystal into the supercooled liquid and begin stirring both test tubes (gently of course).
NOTE:
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Do not attempt to move the thermometer when it is frozen in the sodium thiosulfate
pentahydrate. As the material solidifies, you may not be able to stir very much.
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8. When the temperature of the liquid in test tube #1 (non-supercooled liquid) falls to 51°C, drop in
a seed crystal of sodium thiosulfate to"seed" the liquid, thus aiding crystallization and preventing
supercooling. Continue stirring (gently of course).
9. Continue collecting data for the time listed on the graph (30 minutes); until the temperature
drops to ~ 40°C in both test tubes, or until the cooling curves level out and remain about level for
5 minutes
10. Stop the collection process.
•
Print your graph of the data and attach it to the end of the experiment. (File Print Graph)
•
Be sure to label the graph and indicate the identity of each curve. (by hand is ok)
•
Analyze Data A. To determine the freezing temperature of the sodium thiosulfate
pentahydrate, you need to determine the mean (or average) temperature in the portion of
graph with nearly constant temperature. Move the mouse pointer to the beginning of the
graph’s flat part for the non-supercoooled substance. Press the mouse button and hold it
down as you drag across the flat part of the curve, selecting only the points in the plateau.
Click on the Statistics button,
. Select the correct temperature probe for the nonsupercooled curve. The mean temperature value for the selected data is listed in the statistics
box on the graph. Record this value as the freezing temperature of the non-supercoooled
sodium thiosulfate pentahydrate below. Click on the upper-left corner of the statistics box to
remove it from the graph.
Analyze Data B. Repeat this process for the supercoooled sodium thiosulfate pentahydrate,
Record below.
•
11. When you are finished, heat the solids gently to remove the temperature probes. The solid
should be left in the test tube and returned to the reagent table to be used by the next group.
Rinse the temperature probes with water and dry them carefully to remove any remaining solid.
Questions:
1. Melting/freezing point the of non-supercooled liquid is ___________°C
2. Melting/freezing point of the supercooled liquid is ___________°C
3. Are the freezing/melting points of the non-supercooled and supercooled liquids approximately
the same? Comment on any differences between the two values.
4. When the seed crystal was added to the liquid, record your visual observations of what happened
to the solution in the test tube:
5. A. Did the temperature of the supercooled liquid rise or fall upon addition of a seed crystal?
B. Was this an exo- or endo-thermic process? ________________________
6. What would the final temperature of both test tubes be if we let them sit for 8 hours?
_____________________
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pre-lab stamp here
Intermolecular Forces and Solid-Liquid Transitions
1. Complete the table:
Substance
Formula
Structural Formula
(condensed structural formula is ok)
ethanol
C2H5OH
1-propanol
C3H7OH
1-butanol
C4H9OH
pentane
C5H12
methanol
CH3OH
hexane
C6H14
Molecular
Weight (g/mol)
Hydrogen Bond
(Yes or No)
2. Which will evaporate faster, 1-butanol or pentane? Explain your choice.
3. State the following as either exo- or endo-thermic processes:
Melting _______________________
Freezing__________________________
Vaporization ___________________
Condensation ______________________
Sublimation ____________________
Deposition _______________________
turn over for the last pre-lab question . . .
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4. A thermometer is placed in a test tube of chipped ice at -5.0°C. The temperature is recorded at
the time intervals shown below until room temperature is reached. Plot the data on graph paper
(or using a graphing program like excel). Include a title, labeled axes, and write an explanation
for what phases are present/what is occurring for both flat portions of the curve. Please write the
explanation ON the plot above the horizontal regions. Plot time on the x-axis!
Time (min)
0
2
4
6
10
15
20
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Temp (°C)
-5.0
-2.5
-1.0
0.0
0.0
0.0
0.0
Time (min)
25
30
35
40
45
Temp (°C)
0.0
1.5
4.0
8.0
11.5
50
15.0
55
17.5
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Time (min)
60
65
70
75
80
Temp (°C)
19.0
20.0
20.0
20.0
20.0
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