To determine the enthalpy of vaporisation of water from the measurement of vapour pressure at various temperatures.
The Clausius - Clapeyron equation gives the relationship between the vapour pressure of
a pure liquid and its temperature.
The relationship can be written simply as:
Delta H vap
ln p = - ------------ + constant
R T
where D H is the enthalpy of
vaporisation of the liquid.
The equation is derived from consideration that the Gibbs free energies for a liquid and
vapour are equal when they are in equilibrium.
Several assumptions are involved, the main ones are:
(i) the volume of vapour is assumed to be much greater than that of the liquid vapourised,
and
(ii)it is assumed that the vapour behaves like an ideal gas.
In this experiment, a sample of air is trapped over water, in an inverted measuring
cylinder in a beaker. When the temperature of the apparatus is changed the number of moles
of water vapour in the gas phase will vary according to the Clausius-Clapeyron equation,
while that of air will remain constant.
The number of moles of air in the mixture can be found by reducing the temperature of the
whole apparatus to about 5° C. At that temperature it can be assumed that the
vapour pressure of water is so small that the volume of gas measured corresponds only to
the air present.
The enthalpy of vaporisation can then be calculated from a plot of ln p(H2O)
(the vapour pressure) versus 1/T.
10 cm cylinder, thermometer (preferably one reading to ~ 0.1° C), a tall beaker, bunsen burner.
1. Fill a 10-cm3 graduated cylinder about 80% full with distilled water. Cover the top with a finger and quickly invert and lower the cylinder into a tall beaker that has been filled with tap water. An air sample of 3 to 4 cm3 should be trapped within the cylinder, record this volume and the temperature.
2. Add more water if necessary to the beaker to ensure that the
whole cylinder is surrounded by water. Then heat with a
Bunsen
burner to approximately 80° C. During the heating,
record the time, the
volume and the temperature at every 5° C.
3. When the volume of trapped gas expands beyond the scale on the
cylinder, remove the burner and allow the water to cool slowly.
When the gas begins to contract and the volume can be read again,
record the volume and temperature to the closest 0.1 cm3
and
0.5° C respectively. Stir the water bath frequently
to avoid
thermal gradients. As the water cools, make additional T
measurements at approximately 0.2 cm3 intervals down
to 50° C.
You should be able to record at least 15 readings.
4. After the temperature has reached 50° C, cool the water
rapidly to less than 5° C by adding ice. Record the air
volume
and the water temperature at 10 mins after reaching ca 5° C.
By then an equilibrium has been reached again.
5. Obtain a value of the atmospheric pressure from the Demonstrator.
1. Correct all volume readings by subtracting 0.2 cm3 to
compensate
for the inverted meniscus. Using the measured values for
volume and
temperature from step 4 and the atmospheric pressure, calculate
the
number of moles n(air) of trapped air. Assume that the
vapour pressure of
water is negligible compared to atmospheric pressure at the low
temperature.
2. For each temperature between 80 and 50° C calculate the partial pressure of air in the gas mixture.
n(air) RT
p(air) = ------------
V
3. Calculate the vapour pressure of water at each temperature:
p (H2O) = p (atm) - p (air)
4. Plot ln p(H2O) versus 1/T and draw the best straight line.
5. Determine delta H(vap) from the slope and p(H2O) at the
temperature, X,
given in class. Determine the standard deviation of delta H
(vap) by the
"box method".
(i) Are there any assumptions other than those already mentioned
which are worth considering?
(ii) Discuss the main sources of errors.
Return to Chemistry, UWI-Mona, Home Page
Created March 1995. Last modified 12th April-98.
URL http://wwwchem.uwimona.edu.jm:1104/lab_manuals/c10p5.html