VAPOR PRESSURE
VAPORIZATION
AND
THE
ENTHALPY
OF
INTRODUCTION
When a volatile liquid is placed into a container it starts to evaporate. If the container is left open to the
atmosphere all of the liquid will evaporate. However, if the container is then closed only a portion of the
liquid will evaporate. At a given temperature the rate at which the molecular liquid evaporates is constant.
As more of the molecules are in the vapor phase some of them start to re-enter the liquid phase. Initially,
the rate at which this occurs is low due to the low number of molecules in the vapor phase. Eventually,
there are enough molecules in the vapor phase such that the rate at which the molecules re-enter the liquid
phase and the rate at which they evaporate are equal to each other. The partial pressure of the volatile
substance in the container is the vapor pressure of that substance at that temperature.
If we increase the temperature the vapor pressure will increase and vice versa. The relationship between
the vapor pressure and the temperature is
∆ H
 P 
ln    =
− vap + C
RT
P 
where P is the vapor pressure, P is the unit pressure in the same units as P, ΔvapH is the enthalpy of
vaporization, R is the ideal gas constant in energy units, T is the absolute temperature and C is a constant
that depends on the substance. If we measure the vapor pressure of a substance at various temperatures
and plot the natural logarithm of the vapor pressure on the y-axis and the inverse of the absolute
temperature on the x-axis we should get a straight line with a slope equal to − ∆ vap H . From the slope of the
plot we can then determine the enthalpy of vaporization.
R
PROCEDURE
1.
Put about 250 mL of water into a 400-mL beaker (from your locker) and place it on a hot plate. Set
the hot plate to about 5.
2. Go to one of the experimental setups in the lab. Each set up is comprised of a Vernier LabQuest
device, a Temperature Probe and a Gas Pressure Sensor with syringe and accessories.
3. Connect the Temperature Probe to CH1 on the top of the LabQuest and
connect the Gas Pressure sensor to CH2.
4. Turn on the LabQuest by pressing the power button at the top left corner
of the front of the LabQuest.
The display should now show the
temperature in °C in the box on the top and the pressure in kPa in the box
on the bottom.
5.
Place the white stopper from the Gas Pressure Sensor into a 125-mL
Erlenmeyer flask (from your locker). Connect the white tube on the stopper
to the Gas Pressure Sensor with the clear tubing provided. Leave the valve
on the blue side open for now.
6. Place enough water into a 600-mL beaker (from your locker) such that the
water covers the Erlenmeyer flask up to the stopper. Clamp the flask in
place as in Figure 1. Place the Temperature Probe into the water. Allow
the flask and the Temperature Probe to remain in the water until the
Figure 1
temperature reading stabilizes and then close the valve on the blue side of the stopper. Record the
temperature and the pressure of the air in the flask.
7.
Use the syringe provided to obtain 3.0 mL of ethanol. Connect the syringe to
the blue side of the stopper (Figure 2). Open the valve and depress the
plunger to transfer the ethanol to the flask. Immediately pull the plunger back
up to 3.0 mL and close the valve.
8. Allow the system to come to equilibrium. This may take 30 seconds or more.
You may need to stir the water in the beaker with the Temperature Probe to
keep the temperature constant throughout the bath. Record the temperature
and total pressure for Trial 1.
9. Use the pipet and pipet pump to remove approximately 25 mL of water from
the 400-mL beaker.
10. Add some of the water from the hot water bath that was prepared in step 1 to
the 400-mL beaker to raise the temperature by 3 °C to 4 °C.
Figure 2
11. Stir the water in the 400-mL beaker and allow the system to come to equilibrium again. Record the
temperature and the total pressure in the flask for Trial 2.
12. Repeat steps 9 through 11 until you have completed 5 trials.
13. Because the pressure of the air in the flask increases with the increase in temperature we need to
calculate what the pressure of the air is in the flask at each of the temperatures for Trials 1 through
5. The pressure is related to the temperature (at constant volume and number of moles) through
Boyle’s Law.
14. The vapor pressure of the ethanol is calculated by remembering that the total pressure in the flask
is the sum of the pressure of the air in the flask and the vapor pressure of the ethanol.
15. Calculate the natural logarithm of the vapor pressures, ln(P/P), where P is 1 kPa.
16. Convert all of your temperature readings to Kelvin.
17. Calculate the inverse of the absolute temperatures.
18. Prepare a graph (either by hand or in Excel) with ln (P/P) plotted on the y-axis and 1/T (K−1)
plotted on the x-axis. Follow all of the rules for properly preparing a scientific graph. Remember
to turn in your graph with your report.
19. Determine the line of best fit for your data. Calculate the slope of this line.
20. From the slope calculate the enthalpy of vaporization.
Report for Vapor Pressure and …
Name ____________________________________
Section ____________________________________
Pressure of air in empty
flask (kPa)
Trial 1
Trial 2
Total
Pressure
(kPa)
Calculated Calculated
Pressure of
air (kPa)
Vapor
Pressure, P
(kPa)
ln (P/P)
Temperature
(°C)
Temperature
(K)
1/T (K−1)
Enthalpy of
vaporization
(kJ mol−1)
(from graph)
Show Calculations (use additional sheets if necessary)
Initial temperature of
water bath (°C)
Trial 3
Trial 4
Trial 5
Calculated
Calculated
Calculated
Questions for Vapor Pressure and …
Name ____________________________________
Section ____________________________________
1.
The accepted value for the enthalpy of vaporization of ethanol is 42.32 kJ mol−1. Calculate the
percent error of your experimental value. Refer to Appendix A regarding how % error is
determined. Comment on the relative size of your percent error. Give reasons for why it is
particularly large or small.
2. A substance has an enthalpy of vaporization of 35.76 kJ mol−1. If the vapor pressure is 62.53
mmHg at 22.5 °C, what is the temperature when the vapor pressure is 597.4 mmHg?
Questions for Vapor Pressure and …
Name _______________________________
Section ______________________________
3. A substance has a standard boiling point of 68.72 °C. The same substance has a vapor pressure
at −25.41 °C of 0.113 bar. What is the enthalpy of vaporization in kJ mol−1? What will the vapor
pressure be at 25.00 °C.
4. At what temperature, in K, is the vapor pressure of water equal to 1.500 atm? Data you may
need can be found at the front of the laboratory manual.