Vapor Pressure of Liquids & Solutions
Revised 12/13/14
VAPOR PRESSURE OF LIQUIDS & SOLUTIONS
Adapted from "Chemistry with Computers"
Vernier Software, Portland OR, 1997
INTRODUCTION
When a liquid is added to a closed system (a sealed Erlenmeyer flask, Figure 1), a net flow of
molecules into the gas phase will occur until the rate of evaporation is equal to the rate of
condensation. At this point, the vapor pressure of the liquid is equal to the partial pressure of its
vapor in the flask.
Figure 1
In this experiment, pure ethanol or a solution containing ethanol will be injected into a sealed
Erlenmeyer flask. A computer-interfaced pressure sensor will be used to measure changes in the
total pressure in the flask as temperature of the surrounding water bath is increased. By plotting
the natural log of the vapor pressure (ln P) versus inverse Kelvin temperature (1/T), the heat of
vaporization (ΔHvap) will be determined from the slope of the best fit line using the Clausius
Clapeyron equation (eqn 1):
(1) ln P = (–ΔHvap / R) (1/T) + C
(R = 8.3145 J/mol·K & C = constant)
A colligative property of solutions is the lowering of the vapor pressure of a solvent in solution
in comparison to a solvent by itself. Solutions containing nonvolatile solutes follow Raoult’s
Law (eqn 2).
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(2) Psolution = P°solvent Χsolvent
(P°solvent = vapor pressure of the pure solvent,
Χsolvent = mole fraction of solvent in solution)
Solutions that follow Raoult’s Law are ideal. Intermolecular forces between solvent and solute
in such solutions are very weak. Solution formation results in a negligible change in enthalpy
(ΔHsolution ~ 0). The vapor pressure of nonideal solutions deviates from Raoult’s Law. The
direction of the deviation is dependent on the type of intermolecular forces present between
solute and solvent. If attractive forces are present the deviation will be negative: the actual
vapor pressure will be less than that predicted by Raoult’s Law. Solution formation is
exothermic (ΔHsolution is large & negative). If repulsive forces are present the deviation will be
positive: the actual vapor pressure will be greater than that predicted by Raoult’s Law. Solution
formation is endothermic (ΔHsolution is large & positive).
SAFETY
Wear safety goggles and lab apron at all times in lab. Ethanol and acetone are flammable; make
sure there are no sparks or flames in lab. Ethanol and acetone fumes can irritate eyes and lungs;
avoid contact with fumes. Be sure to transfer all data to your ELN.
Before starting the experiment, the TA will asks you to do a quick demonstration or talk-through
one of the following:
1) Assemble the glassware setup for this experiment (without the parafilm)
2) How to use parafilm? Parafilm the glassware setup assembled in the previous step.
Make sure you watch the videos on the course website and read the documents to prepare. These
demonstrations will be done every week. Everyone will have presented at least one topic by the
end of the quarter. The demonstrations should be short (>1 min) and will be graded.
PROCEDURES
Part A. Vapor Pressure of a Pure Liquid
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1. Work in pairs. Wear safety goggles and lab apron at all times in lab. Calculate air and vapor
pressure as you collect each pressure, temperature data pair (see Calculations #1) – many
students get erroneous results for this experiment, so calculating as you go will enable you to
repeat measurements. Do not inhale fumes and wash hands thoroughly before leaving lab.
Ethanol is very flammable, extinguish any open flames in lab.
2. Heat about 400 mL of tap water in a beaker to boiling temperature on a hot plate.
3. Prepare the glassware setup for the experiment.
•
Obtain a rubber-stopper assembly with a piece of heavy-wall plastic tubing connected to
one of its two valves (Figure 2).
Figure 2
•
Insert the rubber-stopper assembly into a 125-mL Erlenmeyer flask. Important: Twist
the stopper into the neck of the flask to insure a tight fit. Attach the connector at the free
end of the plastic tubing to the gas pressure sensor valve with a clockwise turn. Wrap
parafilm around the stopper and top of the flask to provide extra insurance against gas
leaks. (Parafilm should be stretched when doing this or it will not provide a good seal.)
c. Close the 2-way valve above the rubber stopper as shown in Figure 3. Do this by turning
the white valve handle so it is perpendicular with the valve step itself.
Figure 3
4. Obtain 3.0 mL ethanol (EtOH) using a syringe. With the two-way valve still closed, screw
the syringe onto the two-way valve as shown in Figure 4, but do not place the flask into the
water bath until directed to do so.
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Vapor Pressure of Liquids & Solutions
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Figure 4
5. Plug the temperature probe into CH 1 and the pressure sensor into CH 2 of the LabQuest2
interface box. Make the following adjustments:
•
Go to the
screen, click on graph from the menu and then graph options… change X-
axis column to “temperature” and under graph 2  Y-axis, unselect CH1:Temperature.
Click OK.
•
Go to the
screen, click on mode and change to “selected events.” Click OK.
•
Change the units on the pressure sensor to mmHg. Do this by clicking on the box labeled
“CH2:Pressure” and in the drop down menu select Change units  mmHg.
6. Take a reading of the pressure of air trapped in the Erlenmeyer flask. Read the air pressure
in mm Hg from the
screen. The pressure reading of air in the flask should be the same as
atmospheric pressure (which is around 760 mm Hg on a mild sunny day). Suspend the temp
probe in the air (away from the hot plate!) and read the room air temperature in °C from the
live display below the graph. Record these values in a data table.
7. Inject the ethanol into the flask:
•
Open the 2-way valve above the rubber stopper: turn
the white valve handle so it is aligned with the valve
stem itself as shown here.
•
Inject the EtOH into the flask by pushing in the
plunger of the syringe.
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Vapor Pressure of Liquids & Solutions

Revised 12/13/14
Quickly close the 2-way valve by turning the white valve handle so it is perpendicular
with the valve stem.

Remove the syringe from the 2-way valve with a counter-clockwise turn taking care not
to loosen the stopper assembly from the flask. Leaks will give inaccurate data in this
experiment!
8. To monitor and collect pressure and temperature data:
•
Click
in the
screen.
•
After a few moments equilibrium will be established between EtOH liquid and vapor at a
given temperature. (This is indicated when the P and T readings displayed on the
computer monitor stabilize.) Once this happens, click Keep. The first P-T data
measurement (at room temperature) is now stored. (The T measurement should be the
same as the initial T reading in step 8.)
•
Calculate the vapor pressure at room temperature.
9. Place the sealed flask in a room temperature water bath (1 L beaker) with the entire flask
covered as shown in Figure 4 on the previous page. (The flask will be buoyant and may need
to be clamped down.) Place the temperature probe in the water bath. Next, carefully increase
the temperature of the water bath about 3°C by adding a small amount of the boiling water
prepared on the hot plate. (Use a pipet to measure out hot water or pour the hot water from
the beaker protecting your hands with paper towels.) Stir the water bath with the temperature
probe. Make sure the flask is still covered with water, wait until P and T readings on the
computer monitor stabilize and then click Keep. Your second PT data pair is now stored.
Calculate the air and vapor pressure at this temperature.
10. Repeat the step above by adding small increments of very hot water and collecting PT data.
Calculate the air and vapor pressure at each temperature. (Because the capacity of the beaker
is only one liter, you may have to pour off some water from the water bath before adding
more hot water.) Obtain about 6 measurements between room temperature and 40°C. (EtOH
boils at 78°C and above 40°C the stopper begins to pop out of the flask.)
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11. When you have completed your measurements for EtOH, open the 2-way valve above the
rubber stopper to release the pressure inside the flask. Make sure to point the open valve
away from you and lab partner. Remove the stopper assembly and fill the flask to the top
with water and pour the solution in the sink. (If an aqueous solution is 40% or less ethanol it
poses no environmental threat and can be poured down the drain.)
Part B. Vapor Pressure of a Solution
You are required to measure the vapor pressure of all 3 solutions listed below and share your
data with the entire class for data analysis. Provide detailed procedures in your ELN.
Notes:
(1) The water bath is not necessary.
(2) The same Erlenmeyer flask from Part A should be used in Part B. Create the solutions
inside this flask by adding nonvolatile solutes (ethylene glycol, benzoic or lauric acid) then
capping the flask and adding ethanol and volatile solutes by syringe. Make sure to “swirl”
flask with benzoic or lauric acid until solid is completely dissolved.
Solvent
Solute
2.0 mL ethanol
1.0 mL ethylene glycol (HOCH2CH2OH)
2.9 mL ethanol
0.250 g benzoic acid (C6H5C(O)OH)
2.9 mL ethanol
0.250 g lauric acid (CH3(CH2)10C(O)OH)
When finished, all solutions should be discarded in the collection bottles in the hood.
Then provide the following data on the white board:
•
Air Pressure without solution in the flask (from Part A)
•
Total Pressure of flask with ethanol at RT (from Part A).
•
Exact mass or benzoic acid and lauric acid added.
•
Exact volume of ethylene glycol added
•
Total Pressure of flask with the above solution.
Record all other students’ data before leaving lab.
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Vapor Pressure of Liquids & Solutions
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Make sure to clear your email address and password of the LabQuest2 so others can’t access your
email account. Shutdown the LabQuest2 and not simply put it to sleep. To shutdown the LabQuest2:
press the home key, select System  Shut Down  OK.
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Vapor Pressure of Liquids & Solutions
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CALCULATIONS
Almost all of the following work must be done on Excel spreadsheet, using the functions in the
spreadsheet to perform the calculations. In other words, you can’t use your calculator to
perform the calculations, you must enter the formulas in Excel and have Excel do the
calculation. No credit will be given if the TA can’t find formulas written into the cells. You must
also include header cells with units. If you do not know how to use Excel in this way, you should
search the internet for written or video directions. The Excel spreadsheet must be attached to the
ELN.
Part A. Vapor Pressure of a Pure Liquid.
(1) Find the vapor pressure at each temperature. The total pressure in the flask at temperature,
T, is equal to sum of the pressure of air in the flask and the vapor pressure of the ethanol:
(P total)T = (P air)T + (VP ethanol)T.
Remember, for trials at temperatures other than room temperature, even if no ethanol was
present, the air pressure would increase due to a higher temperature, or decrease due to a
lower temperature (remember those pesky gas laws?). Therefore, the air pressure at room
temperature must be corrected to the air pressure at the temperature of the water bath. Use
the gas-law equation: P2 / T2 = P1 / T1 (T should have Kelvin units).
(2) Construct a Clausius-Clapeyron plot:
a) In Excel, construct a graph of vapor pressure of EtOH vs. Celsius temperature. What is
the general relationship between vapor pressure and temperature?
b) Using Excel to perform the calculation, convert vapor pressure of ethanol to ln VP and
Celsius temperature to inverse Kelvin temperature (1/T). Construct a graph of ln VP vs.
1/T. Calculate ΔHvap from the slope.
c) The true value for ΔHvap of ethanol is 42.3 kJ/mol. Calculate your percent error: % error
= [(true – expt)/ true] x 100.
d) To calculate the air pressure, the volume occupied by the air is assumed constant. Why
is this assumption incorrect? Explain how the vapor pressure calculated and the
resulting Clausius-Clapeyron plot are affected?
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Vapor Pressure of Liquids & Solutions
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e) In the Sapling prelab for this experiment you used Spartan to calculate the enthalpy of
formation (ΔHf) of the hydrogen bond between two ethanol molecules. Compare that
value with the enthalpy of vaporization (ΔHvap) you found experimentally for ethanol.
The magnitudes of the two values should be close, but the sign is opposite. Why? What
are the values and how are they related?
Part B. Vapor Pressure of a Solution.
(1) Find the experimental vapor pressure at room temperature for all measurements. The total
pressure in the flask at temperature, T, is equal to sum of the pressure of air in the flask and
the vapor pressure of the solution:
(P total)T = (P air)T + (VP solution)T. (You should
have ≤12 vapor pressures for each solution type (≤36 vapor pressures). Remember, this
should be done in Excel.)
(2) Use the Q-test to test the lowest and highest values for the vapor presures of each solution
calculated in #1 (You’ll need to do 6 Q-Test, 2 for each solution.) Reject any value that falls
below the 90% confidence level. (Do not use the measurements in the calculations that
follow.)
(3) Using Raoult’s Law, calculate the theoretical vapor pressure at room temperature for each
solution measurement. The vapor pressure for pure ethanol at room temperature is 58.71
mmHg.
(4) For each solution type, find the mean vapor pressure and the relative average standard
deviation for the both the experimental and theoretical vapor pressure. (Again, you can use
Excel to this very quickly. You should have 6 mean vapor pressures with 6 relative average
standard deviations accompanying them.)
(5) Determine the percent error between the experimental and theoretical values calculated in
#4. Do NOT use absolute values in the calculation. Indicate if the solution is ideal or
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Vapor Pressure of Liquids & Solutions
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nonideal – if nonideal indicate the direction of deviation from Raoult’s Law: positive or
negative.
QUALITATIVE ERROR ANALYSIS
1. What modifications could be made to the procedure to better account for random
(indeterminate) errors?
2. List three potential systematic (instrumental, methodological, or personal) errors that could
be made in this experiment. (Note: Be specific, systematic errors are in the details. For
example, losing your solution because you knocked over the cuvette is not a systematic error
– it’s a gross one.)
3. Did any gross errors occur? Did you mess up? Did the equipment or instrumentation fail?
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