Experiment 12Z
(Adapted from Expt 12E by MIDN 1/C Cabarrus & 1/C Brown)
FV 4-18-16
INTERMOLECULAR FORCES AND THE LIQUID-VAPOR
EQUILIBRIUM1
MATERIALS: 150 mL beaker, 6 mL graduated plastic syringe sealed at the tip, digital thermometer, hot
plate, plastic bin, 2” metal washers
PURPOSE: The purpose of this experiment is to measure the effect of temperature on the vapor pressure of
several liquids. The data will be analyzed to extract values for the heat of vaporization, ΔHvap, and interpreted
in terms of intermolecular forces.
LEARNING OBJECTIVES: By the end of this experiment, the student should be able to demonstrate the
following proficiencies:
1.
2.
3.
4.
5.
Use a spreadsheet program for data manipulation, graphing, and regression analysis.
Describe the effects of changes in temperature on the vapor pressure of a pure substance.
Describe how intermolecular forces influence the relative vapor pressure of a pure substance.
Understand the use of graphical methods to extract thermodynamic information from experimental
pressure and temperature data.
Utilize Dalton’s Law of Partial Pressures, and the Ideal Gas Law, to relate experimental data to
properties of the test substance.
DISCUSSION:
The molecules of a gas move freely throughout the entire volume of the container, the individual molecules
staying widely separated and experiencing little or no interaction with other molecules. Molecules in a liquid,
while free to move throughout the volume of the sample, are constrained by intermolecular forces to remain
in contact with their neighbors. The strength of such intermolecular forces and the energy of motion available
to the sample (based on the temperature), together dictate the physical state of a substance.
Evaporation is the process of converting a substance from the liquid phase to the gas phase. It is an
endothermic process, since energy is required to overcome the attraction that a liquid molecule feels for its
neighbors. The molar enthalpy of vaporization, ΔHvap, is the energy required to evaporate one mole of a
substance at constant temperature and pressure. The magnitude of ΔHvap is thus a measure of the strengths
of the intermolecular forces in a pure substance.
The molecules in a liquid will have a distribution of energies at any temperature, as do the molecules of a
gas. If a liquid is placed in an evacuated, closed container, some of the molecules of the liquid (those in the
higher energy range) will have sufficient energy to escape to the gas phase. Thus the pressure in the container
will rise. Some of the gas phase molecules will hit the liquid surface and be unable to escape the attractions
for their new neighbors; these (lower energy molecules) have undergone condensation and become part of
the liquid. As more molecules accumulate in the gas phase (via evaporation), the rate of condensation will
also increase. Eventually, the rate of evaporation and the rate of condensation will become equal, and the
pressure in the container will level off at some constant value. The system is said to be in equilibrium, and
the pressure of gas that exists over the liquid is called the equilibrium vapor pressure of the liquid.
The vapor pressure depends on the temperature of the sample (since a higher temperature gives a larger
fraction of high-energy molecules), and on the strength of the intermolecular forces holding molecules of the
1
Levinson, G.S., Journal of Chemical Education, 59, 337 (1982).
1
liquid together. The latter connection allows a determination of the enthalpy of vaporization by a study of
the vapor pressure (P) as a function of temperature:
(1) where R is the gas constant (8.314 J/mole·K), T is the absolute temperature (in K), and C is a constant. As
seen in this equation, liquids with a large positive value of ΔHvap will have a low equilibrium vapor pressure
at any temperature. As the temperature increases, ln P, and thus P, also increases.
In this experiment, the volume of a gas mixture of air and sample vapor will be measured at several
temperatures and at atmospheric pressure. One such measurement will be made near the freezing point of
water, when the vapor pressure of water or the sample organic compound is nearly zero. This allows a
determination of the (constant) number of moles of air trapped in the cylinder and thus the partial pressure
of air in the mixture at any temperature. From this and the barometric pressure, the partial pressure of the
liquid at any temperature can be determined by their differences. Then the application of Eq. (1) allows the
determination of the molar heat of vaporization ΔHvap by graphical methods. 2
PROCEDURE:
1.
Obtain the barometric pressure and record it in the attached data sheet. Take the small plastic bin
from your student drawer and fill it with ice.
2.
Fill a 150 mL beaker with distilled water. Place about 6 mL of distilled water in the prepared
syringe. Carefully cover the top of the syringe with your finger, invert it, and place it in the beaker.
Do not release the syringe until you have placed a metal washer atop the syringe to hold it under
the water. The syringe must be completely immersed throughout the experiment.
3.
ORGANIC STEP: See your Instructor for a sample loading
syringe containing your assigned organic liquid. (If assigned
water as a sample, just continue to step 4.) You or your
Instructor should load the sample by carefully placing the
opening of the injector tube underneath the opening of the
submerged syringe. Slowly inject the 0.5 mL of organic,
watching it rise to the top of the submerged syringe.
4.
Record the identity of your sample liquid on the data sheet.
Place the newly prepared apparatus on the heating mantle and turn the heat dial to the highest
setting.
5.
Using a thermometer, continuously stir the water while it heats. This is done to ensure that the
temperature of the water bath is uniform throughout.
6.
While continuing to stir, diligently monitor the level of the gas bubble within the submerged
syringe. As the temperature rises, the gas bubble will expand. Should it begin to expand too
quickly, remove the beaker from heat, wait about a minute, and begin heating again. (If any of the
gas bubble escapes, you will have to start over - see your Instructor.)
7.
As the level of the gas bubble reaches the 6 mL mark, remove the beaker from heat and allow to
cool. Continue to stir throughout the cooling process.
8.
Once the gas bubble reaches the 4 mL mark on the syringe, record the temperature and volume at
every 0.2 mL of volume change. Record this data in the attached data sheet to the nearest 0.1oC
and 0.1 mL. Continue to stir throughout the measurement process. Make sure the tip of the
thermometer is in the middle of the beaker, not at the bottom, when you measure temperature.
9.
After 11 data points have been collected, cool the beaker rapidly to below 5ºC. To do this, place
your 150 mL beaker and syringe in the plastic bin of ice, without losing the air trapped in the
syringe. Continue stirring and allow the bath to cool completely before recording the temperature
and gas bubble volume. Record the final cooled temperature and volume in the attached data
sheet.
Clean-up:
1.
Remove the 150 mL beaker plus syringe from the ice container. Remove your syringe from the
150 mL beaker, empty any liquid remaining into the beaker, and place the syringe in the container
labeled for the organic sample it contained.
2.
Empty the water from your 150 mL beaker into the waste container in the Instructor’s hood, NOT
IN THE SINKS. The water is contaminated with a small amount of organic liquid and must be
treated separately.
3
Section _________________________
Name _________________________
Date _________________________
Lab Partner ________________________
DATA SECTION
Experiment 12Z
Sample Liquid ______________________
Barometric Pressure (mm Hg) ____________
Gas Bubble Volume (mL)
Temperature (ºC)
Gas Bubble Volume (mL)
Temperature (ºC)
After the addition of ice:
4
DATA TREATMENT
Experiment 12Z
1. Set up an Excel spreadsheet, marking columns for the Volume (mL) and Temperature data (°C), as well
as columns for Volume (L), Absolute Temperature (K), Moles of Air, Partial Pressure of Air (atm), Partial
Pressure of Sample, Inverse Kelvin Temperature and Ln(Psubstance) values. Insert your volume and
temperature data into the spreadsheet.
2. Convert all volumes into Liters. Convert all Celsius temperatures into absolute (Kelvin) temperatures.
Convert barometric pressure to Atmospheres (atm). Use these volumes and absolute temperatures in all
remaining calculations.
3. Determine the number of moles of air, nair, trapped in the gas bubble using your gas volume after the
addition of ice and the barometric pressure. Treat the gas bubble as an ideal gas: nair = PbarV/RT. This assumes
that the pressure of gas in the bubble is the same as the atmospheric pressure in the room. At the low
temperature used for this calculation, the vapor pressure of water or organic sample is negligible. (Check a
table to see if this is a good assumption!) Thus, the gas bubble consists essentially of air alone at this
temperature. The number of moles of air does not change throughout the experiment, so we will be able to
determine how much of the gas bubble is air (and how much is the test substance) at any temperature.
Show your calculation of nair.
4. At all of the high temperature data points, the gas bubble is a mixture of air plus sample vapor. For each
of the high temperature data pairs, calculate the partial pressure of air in the gas bubble: Pair = nairRT/V,
recognizing that the number of moles of air does not change from step 3 above. Show a sample calculation
for your highest temperature data point.
5. For each of the high temperature readings, calculate the partial pressure of the test substance, Psubstance in
the gas bubble: Psubstance = Pbar - Pair. Show a sample calculation for your highest temperature data point.
5
6. Plot Psubstance vs. T for all of the high temperature points (i.e. do NOT include the data from the ice bath). 7. Plot Ln(Psubstance) vs. 1/T for the same points. Perform a linear regression (trendline) analysis on the latter,
and plot the best straight line.
Equation of trendline: _________________________________
Slope (with units): ____________
R2 for trendline: _____________
8. According to equation (1), the slope of your plot is related to the heat of vaporization of the sample. From
the slope of your regression line, determine the molar heat of vaporization, ΔHvap, and compare it to the
accepted value of your test substance. Calculate the percent deviation. Include these results on your
spreadsheet printout.
Literature Values:2
Heat of Vaporization
(∆Hvap)
40.65 kJ/mol
31.5 kJ/mol
36.1 kJ/mol
46.5 kJ/mol
Solution
Water
Hexane
Heptane
Nonane
Molar Mass
(g/mol)
18.0153
86.1754
100.2019
128.2551
ΔHvap (show work):
Percent Deviation (show work):
2
Glushko Thermocenter, “Entropy and Heat Capacity of Organic Compounds” in NIST Chemistry WebBook,NIST Standard
Reference Database Number 69, Eds. P.J. Linstrom and W.G. Mallard, National Institute of Standards and Technology,
Gaithersburg MD, 20899, http://webbook.nist.gov, (retrieved April 8, 2016).
6
9. Enter your result in the Class Data table. Record data for the other samples from the Class Data table, so
that you have experimental ΔHvap values for all samples used.
CLASS DATA
Experiment 12Z
Sample
Heat of Vaporization (∆Hvap) (kJ/mol)
Average
(kJ/mol)
Water
Heptane
Nonane
Hexane
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QUESTIONS
Experiment 12Z
1.
Rank the liquids based on your experimental ΔHvap values from greatest to least. 2.
Using the structures provided3, circle all of the intermolecular forces present in each pure liquid.
Heptane
Nonane
Hexane Water
Nonane -
Dipole-Dipole
LDF
Hydrogen Bonding
Water -
Dipole-Dipole
LDF
Hydrogen Bonding
Heptane -
Dipole-Dipole
LDF
Hydrogen Bonding
Hexane -
Dipole-Dipole
LDF
Hydrogen Bonding
(LDF = London Dispersion Forces)
3. Continuing using the structures as guidance and your data collected, explain the ranking
established in Question 1. What is the trend for the order in molecules that only have London Dispersion
Forces (LDF’s)?
4.
How do intermolecular forces relate to/affect a liquid’s heat of vaporization, ΔHvap?
3Glushko
Thermocenter, “Entropy and Heat Capacity of Organic Compounds” in NIST Chemistry WebBook,NIST Standard
Reference Database Number 69, Eds. P.J. Linstrom and W.G. Mallard, National Institute of Standards and Technology,
Gaithersburg MD, 20899, http://webbook.nist.gov, (retrieved April 8, 2016).
8
Section _________________________
Name _________________________
Date _______________________
PRE-LAB Questions
Experiment 12Z
1.
Define the following types of intermolecular forces:
London Dispersion Forces (LDF) -
Dipole-Dipole Interactions -
Hydrogen Bonding - 2.
Define Heat of Vaporization (ΔHvap).
3.
What molecule would you expect to have a higher vapor pressure, a molecule with strong
intermolecular forces, or weak intermolecular forces? Explain your answer briefly.
4.
Referencing the table on page 6, does the molar mass (how heavy the molecule is) matter in the
determination of whether a molecule will have a high or low vapor pressure? Explain your answer
briefly.
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