Experiment 14
Intermolecular Forces rev 1/12
GOAL:
We will examine connections between molecular structure, intermolecular forces, and physical properties.
BACKGROUND:
Physical properties such as solubility, melting point, and boiling point are determined by a substance’s
intermolecular forces. Non-polar molecules have the weak intermolecular forces known as London
dispersion forces caused by slight, temporary asymmetries in their electron clouds. Although London
dispersion forces are relatively weak, as the number of electrons in a molecule increases so does the
strength of the London dispersion forces. So, if we are comparing several non-polar molecules to each
other, we expect the largest ones (with the most electrons) to have strongest London dispersion forces.
Polar molecules have London dispersion forces in addition to dipole
forces. Dipole forces are attractions between + and - charges on the
molecules. Dipole forces are generally stronger than London dispersion
forces because the charges are permanent, not temporary, and the
charges are stronger. Both dipole forces and London dispersion forces
figure into the total strength of intermolecular forces for polar
compounds. Suppose we had two molecules with very similar London
dispersion forces but one was polar and one was not. The presence of the
dipole forces will give the polar molecule stronger intermolecular forces
overall.
-
+
Hydrogen bonds are the strongest intermolecular forces found in covalent
compounds. They are formed by the attraction between a hydrogen atom on
O
one molecule and a nonbonding pair of electrons on a second molecule.
H
H
Hydrogen bonds are found only in molecules where a hydrogen atom is
bonded to fluorine, oxygen, or nitrogen. These molecules also have dipole and
O
London dispersion forces, and all three types of forces contribute to the total
H
H
strength of intermolecular forces. The presence of hydrogen bonding makes
the total intermolecular forces very strong and leads to high melting and boiling points. Having multiple
groups capable of hydrogen bonding makes melting points even higher.
Two molecules with exactly the same types of intermolecular forces may still differ from each other in
the strength and impact of these forces due to general molecule shape or the placement of hydrogen
bonding groups which are very particular about geometry. Linear and planar molecules often exhibit
stronger intermolecular forces because their shape allows the molecules to pack closely and maximize the
interactions.
The combined effects of all intermolecular forces determine physical properties. The stronger the
intermolecular forces a compound has, the more energy will be required to overcome those attractions.
Thus, strong intermolecular forces correspond to high boiling points and less volatile compounds, i.e.,
ones that will evaporate less. In Part 1, you will observe several liquids as they evaporate. All liquids
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cool as they evaporate because they must use energy to overcome intermolecular forces and move into the
gas phase. The liquids that show the most cooling have lost the most molecules due to evaporation.
This indicates a compound with weaker intermolecular forces.
In Part 2, you will measure the melting points of several compounds. In Part 3, you’ll also use Spartan
Student to calculate their dipole moments (polarities). In Part 4, you will use physical models to look at
intermolecular forces in water and aqueous solutions.
Table 1: Names and Structures of Organic Compounds Used in this Experiment
Name
Structure
Name
Structure
H
methanol
H
C
ethanol
O
H
H
1-propanol
pentane
H
H
H
H
H
H
C
C
C
H
H
H
1-butanol
O
H
H
H
H
H
H
H
C
C
C
C
C
H
H
H
H
H
H
hexane
H
H
H
C
C
H
H
O
H
H
H
H
H
C
C
C
C
H
H
H
H
O
H
H
H
H
H
H
H
C
C
C
C
C
C
H
H
H
H
H
H
O
H
C
benzoic
acid
HC
C
HC
C
O
H
C
2-hydroxybenzoic acid or
o-salicylic acid
OH
CH
HC
HC
C
H
H
O
H
C
HC
C
C
HO
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naphthalene
C
H
H
C
H
Intermolecular Forces
C
C
C
CH
OH
H
C
C
OH
OH
C
C
H
4-hydroxybenzoic acid or
p-salicylic acid
C
C
H
C
C
C
C
C
H
H
H
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H
PRE-LAB ASSIGNMENT:
Two of the molecules you will study this week are methanol and ethanol. The structures of these
molecules are given on page 15-2. Copy the structures in your notebook. List all the types of
intermolecular forces these molecules have.
HAZARDS:
Several of the liquids that you will be working with are highly flammable, so no flames are allowed in the
lab. While none of the chemicals used in this experiment pose unusual health hazards, your hands may be
exposed to these chemicals. Wear disposable gloves to limit your exposure. Keep the test tubes
containing the liquids closed when not in use. Part 1 is done in the hood. Look up the MSDS for
methanol (http://hazard.com/msds/index.php, or Trexler 464). Record the following information in your
notebook hazards section:
 appearance, odor, and boiling point (see Physical & Chemical Properties)
 a summary of potential health effects (see Hazards Identification)
 a summary of potential effects if released to the environment (see Ecological Information)
LABORATORY DATA AND OBSERVATIONS:
The in-lab portion of this experiment will be done in teams of two students, but the lab reports will be
done separately. This means each student will need a full set of notebook entries. Note the name of your
in-lab partner in your lab notebook. Remember to record both what you do and what you observe in your
notebook. The four parts of this experiment may be done in any order. Your lab instructor will assign
you to start with a given part in order to avoid congestion.
PART 1 PROCEDURE: Evaporation
1.
Your numerical data for Part 1 will be most easily recorded in a table with four columns labeled:
liquid, maximum temperature, minimum temperature, and . You will need one line for each of
the six liquids tested.
2.
Use one of the computers that are set up in a hood. If it has not already been done for you,
prepare the computer for data collection by opening Experiment 13 from the Roanoke
Experiments folder in Logger Pro. On the Graph window, the vertical axis has temperature
scaled from 5 to 30oC. The horizontal axis has time scaled from 0 to 250 seconds.
3.
Wrap Probe 1 and Probe 2 with square pieces of filter paper secured by wire. Roll the filter paper
around the probe tip in the shape of a cylinder. The paper should be even with the probe end.
4.
You will be testing 6 liquids. The four alcohols are methanol, ethanol, 1-propanol, and 1-butanol.
The two alkanes are pentane and hexane. These have been pre-measured into test tubes. You
may do the liquids in any order. Get two liquids at a time. When finished, return the liquids to
the supply bench so that others may use them. Keep the liquids stoppered when they are not in
use.
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5.
Wear gloves for the rest of Part 1. Stand Probe 1 in the first liquid container and Probe 2 in the
second liquid container. Record which probe goes into which liquid. Make sure the containers
do not tip over.
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
. Monitor the temperature for 15 seconds to establish the initial (maximum)
temperature of each liquid. Then simultaneously remove the probes from the liquids and tape
them to the metal flashing at the front edge of the hood so that the probe tips extend 5 cm over the
edge and into the hood. The probe tip should not touch the working surface of the hood. Pull the
hood sash down as far as you can. Stopper the test tubes to minimize the loss of these liquids.
8.
When both temperatures have reached minimums and have begun to increase, click
to
end data collection. Two statistics boxes should appear on the screen. You may need to drag the
boxes apart since the computer sometimes places them on top of each other. Record the
maximum and minimum values for Temperature Probe 1 (first liquid) and Temperature Probe 2
(second liquid) in your notebook.
9.
For each liquid, subtract the minimum temperature from the maximum temperature to determine
T, the temperature change during evaporation. Record this in your table.
10.
Roll the rubber band up the probe shaft and dispose of the filter paper in the waste container in
the hood.
11.
Return your first two liquids to the supply bench. Be sure they are stoppered. Get two different
liquids. When you choose Collect for your second set of liquids, you will be prompted to discard
your prior set of data. Discard it. Repeat the procedure above first for liquids 3 & 4, and then for
liquids 5 & 6.
12.
Be sure that both you and your lab partner have a full set of data before leaving lab. Remember
that you will write up your reports separately.
PART 2 PROCEDURE: Melting points of organic solids
1.
One member of your team should prepare melting point capillaries of benzoic acid and
naphthalene. The other team member should prepare melting point capillaries of 2hydroxybenzoic acid (or o-salicylic acid) and 4-hydroxybenzoic acid (or p-salicylic acid).
Remember that you need just one crystal of each in your capillary. Set each capillary on a
labeled piece of paper so that you don’t mix them up.
2.
Measure the melting point range of each solid, starting with the temperature where you first see
softening to the temperature where the entire sample has melted. Put both capillaries in the
MelTemp at the same time. Make sure you know which is which. You will heat rapidly until
you near the melting point for each compound and then slow the rate of heating to just 1o every
10-15 seconds.
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3.
If you are doing the benzoic acid/naphthalene set, heat quickly to 70o then adjust so that
temperature increases just 1o every 10-15 seconds until your first solid melts. Now heat quickly
to 110o and then again adjust so that temperature increases just 1o every 10-15 seconds until your
second solid melts. For each solid, you should record a melting point range of just a degree or
two in which the solid softened and melted.
4.
If you are doing the 2-hydroxybenzoic acid/4-hydroxybenzoic set, heat quickly to 145o then
adjust so that temperature increases just 1o every 10-15 seconds until your first solid melts. Now
heat quickly to 205o and then again adjust so that temperature increases just 1o every 10-15
seconds until your second solid melts. For each solid, you should record a melting point range of
just a degree or two in which the solid softened and melted.
5.
As soon as your second solid melts, get the apparatus cooling back down using the compressed
air. Dispose of the capillaries in the broken glass box.
PART 3 PROCEDURE: Using Spartan Student to calculate dipole moments of organic solids
1. Use Spartan Student to draw models of the four molecules from Part 2: benzoic acid, 2hydroxybenzoic acid, 4-hydroxybenzoic acid, and naphthalene. Follow the structures which are
shown on page 2. Have these drawings handy to follow. Recall that you used Spartan Student
for Exp 11 in CHEM 111. Begin by opening Spartan Student on one of the lab computers.
2. Open a new file by clicking the
button. On the righthand side of the screen, you should see
a collection of molecular fragments that you can use to build the molecules.
3. Start by building benzoic acid. Add a benzene ring to the main drawing area. Now click on the
Groups dropdown menu. Choose Carboxylic Acid. Add this to an open bond in your drawing.
4. Spartan should recognize your drawing as benzoic acid. Look at the bottom right of the screen.
Click the up arrow next to the name. Choose Replace.
5. Now click View:
Click and drag to rotate the molecule.
6. From the dropdown menu at the top of the screen, choose Display, Properties. Record the
displayed dipole moment in your notebook.
7. You can easily modify this benzoic acid structure into 2-hydroxybenzoic acid, also known as
salicylic acid. Click Add Fragment:
8. Choose the molecular fragment at right that is an oxygen atom with two single bonds.
Add this to the appropriate place on your molecule. Spartan should recognize your
molecule. Follow steps 4-6 above to find your dipole moment.
9. Again modify your structure, this time to form 4-hydroxybenzoic acid. Click on Add
Fragment as in Step 7. Now delete the –OH group from your molecule using the
Delete Button: Switch back to Add Fragment, and add the new –OH group on the
correct spot.
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10. Again, Spartan should recognize your molecule as p-salicylic acid. Follow steps similar
to Steps 4-6 above to find the dipole moment.
11. The last molecule you need to make is naphthalene. It is sufficiently different from the
others that you should close this file and open a new one using the buttons across the top
of the screen. You do not need to save a copy of the current file.
12. Once you have a blank drawing area, choose the Rings dropdown menu and select
Naphthalene. You do not need to modify this structure! Simply follow Steps 4-6 above
to record the dipole moment.
13. Once you have dipole moments for your four compounds, simply use the Close
File button but leave the software open for the next group of students.
PART 4 PROCEDURE: Intermolecular forces in water and aqueous solutions
1. Spread out your Cup of Water models. You should have 12 H2O molecules, one Na1+ ion , one
Cl1- ion, one OH hydroxyl group, and one C2H6 (ethane) molecule. Small magnets in the models
will help us simulate intermolecular forces. Record observations in your notebook as you follow
the instructions below.
2. Select two H2O molecules. How are they most attracted to each other: H to H, O to O, H to O, or
some combination of these? Which two types of intermolecular forces are being represented
here? Record your observations.
3. Choose one H2O molecule to be your central molecule. How many other water molecules can
you attach directly to this one central molecule? Draw a structure similar to the one of water
molecules on page 13-1 but showing how all the molecules attach to the central water molecule.
4. In solid H2O (ice), the water molecules arrange themselves so that each water molecule can
hydrogen bond with as many others as possible. Arrange all 12 of your water molecules in the
most compact structure that also maximizes hydrogen bonding. Use a sketch and written
description to record this structure in your notebook. Describe how much open space is trapped
inside your solid structure.
5. In liquid H2O, the water molecules are constantly forming and breaking intermolecular hydrogen
bonds as the individual molecules move. Take your model of solid ice in your hands and
compact it like you would a snow ball to simulate this. When compared to your solid structure
before, how much open space is now trapped inside the liquid structure?
6. Now let’s consider how solutes interact with water molecules. Find the NaCl. The smaller ion is
Na1+. Note the strength of its attraction for a Cl1- ion. How many Cl1- could fit around a single
Na1+?
7. Examine the NaCl model on the instructor’s bench. Use a sketch and written description to
record this structure in your notebook.
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8. How are water molecules attracted to the Na1+: through the O or through the H? Fit the
maximum number of water molecules on the Na1+ possible. Use a sketch and written description
to record this structure in your notebook.
9. How are water molecules attracted to the Cl1-: through the O or through the H? Fit the maximum
number of water molecules on the Cl1- possible. Use a sketch and written description to record
this structure in your notebook.
10. You should now have a Na1+ surrounded by water molecules and a Cl1- surrounded by water
molecules. This is how they actually exit in water. We say that each ion is surrounded by a
“hydration sphere.” Hold the hydrated Na1+ next to the hydrated Cl1-. With the hydration spheres
intact, will the ions stick together?
11. Pull the water off the ions and examine the attraction between the water molecules and the ethane
(C2H6) molecule. How strong is its attraction for water? Does it form a hydration sphere?
12. One of the hydrogen atoms on the ethane molecule is marked with a small colored spot and small
raised dots on the carbon atom at its base. Carefully remove this hydrogen and set it aside in a
safe place. Insert the OH, hydroxyl group, in its place. You now have C2H5OH, ethanol. How
strong is its attraction for water? Use a sketch and written description to record the
structure/interaction between ethanol and water.
13. Remove the OH from your ethanol and replace the H that you previously set aside. Carefully
count the models as you return them to the original cup. Be sure that you have 12 H2O
molecules, one Na1+ ion, one Cl1- ion, and one OH hydroxyl group, and one C2H6 ethane
molecule.
RESULTS:
For Part 1, prepare a Results table with 6 columns: liquid name, formula, structure, molar mass,
intermolecular forces, and T. Structures for the liquids are given in the Introduction. List all the
intermolecular forces expected for each compound. The Introduction or your textbook will provide help.
Combine Parts 2 and 3, preparing a Results table with 7 columns: compound name, structure, molar
mass, experimental melting point, literature melting point, dipole moment, and intermolecular forces.
Structures for the compounds are given in the Introduction (draw them in by hand). Experimental melting
points are the ones you found in lab. Literature melting points are accepted values that you find in
chemical literature. Look up your compounds in either the CRC (see the Table of Physical Properties of
Organic Compounds) or the Merck Index. List the source of your data just below your Results table. In
the final column, list all the intermolecular forces expected for each compound.
QUESTIONS:
Questions 1-4 refer to Part 1 only:
1. Which of the alcohols studied in Part 1 has the strongest intermolecular forces of attraction? The
weakest intermolecular forces? Explain how the results of this experiment show this. Explain how this
result could be predicted from looking at the structures or formulas.
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2. Which of the alkanes studied in Part 1 has the stronger intermolecular forces of attraction? The
weaker intermolecular forces? Explain how the results of this experiment show this. Explain how this
result could be predicted from looking at the structures or formulas.
3. Two of the liquids, pentane (an alkane) and 1-butanol (an alcohol), have nearly the same molar
masses, but significantly different t values. Explain the difference in t values of these substances, based
on their intermolecular forces.
4. Diethyl ether (CH3CH2OCH2CH3) has a molar mass similar to that of pentane and 1-butanol (see
Question 3). Draw the Lewis structure of diethyl ether. Is it polar or non-polar? What types of
intermolecular forces are expected in diethyl ether? What t do you predict for diethyl ether? (Think
about the t you measured for pentane and 1-butanol) Explain your reasoning.
Questions 5-9 refer to Parts 2 and 3 only:
5. How well do your experimental melting points correspond to the literature melting points? What
might cause an experimental melting point to deviate from an accepted value?
6. Rank your four compounds in order of increasing molar mass. Do your melting points correlate well
with molar mass (that is, does increasing molar mass always cause a similarly sized increase in melting
point)? Cite examples from your data to support your statement.
7. Rank your four compounds in order of increasing dipole moment. Do your melting points correlate
well with dipole moment (that is, does increasing dipole moment always cause a similarly sized increase
in melting point)? Cite examples from your data to support your statement.
8. What general connection should exist between melting points and the strength of intermolecular
forces? Give a general explanation of why this is true. (Talk about changes that happen when a solid
melts.) See the Introduction or your textbook for help.
9. Polar molecules should have a measurable dipole moment. Do the dipole moments that Spartan
calculated correspond to your polarity predictions from just looking at the structures? Cite examples from
your data to support your statement.
Questions 10-13 refer to Part 4 only:
10. List all the types of intermolecular forces present in a sample of pure water. Describe each in a
sentence or two. Which of these were simulated by our models?
11. Use what you observed about the structures of solid and liquid water to explain why ice floats on
liquid water. (Hint: think about the effect of that open space you noticed)
12. NaCl dissolves well in water. Describe what you observed with the models to explain why and how
it dissolves.
13. Describe the interactions between water and the two related molecules of ethane and ethanol. Which
will be more soluble in water? Why?
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