Chem1811 Lab talk: Enthalpy of vapourization by vapour pressure measurements Part 1

Main Ideas

Vapor Pressure

Vapor pressure is the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phase (a liquid or solid), usually in a closed or nearly closed system.

Evaporation and condensation occur simultaneously at a liquid’s surface. At equilibrium, their rates are equal. At a given temperature, a liquid’s equilibrium vapor pressure does not depend on the container’s size or the liquid’s exposed surface area. The pressure is produced by vapor molecules moving and colliding with the container.

Evaporation and Boiling

Evaporation and boiling are different processes:

• Evaporation occurs at the surface and can happen at any temperature.
• Boiling occurs throughout the liquid when its vapor pressure reaches or exceeds the external pressure. Vapor bubbles can then form within the liquid.

Because atmospheric pressure is lower at high elevations, liquids boil at lower temperatures there. For example, water may boil at a temperature too low to cook food efficiently at high altitude.

Temperature and Vapor Pressure

As temperature increases, more molecules have enough energy to escape from the liquid into the gas phase, so vapor pressure rises. This increase is nonlinear and becomes more rapid at higher temperatures.

Vapor Pressure and Enthalpy of Vaporization

The Clausius–Clapeyron relationship connects vapor pressure and temperature. Under the usual approximations, plotting (\ln(P_{\text{vap}})) against (1/T) produces a straight line with a slope of

[ -\frac{\Delta H_{\text{vap}}}{R}. ]

The slope can therefore be used to determine the enthalpy of vaporization.

This relationship relies on several approximations:

• The vapor behaves approximately as an ideal gas.
• The vapor’s volume is much larger than the volume of the condensed phase.
• The enthalpy of vaporization changes little over the temperature range studied.

Experimental Approach

The experiment estimates the liquid’s vapor pressure at several temperatures, then uses the pressure–temperature relationship to determine its enthalpy of vaporization.

1. Trap a fixed amount of air in an inverted measuring cylinder over the liquid. The gas inside contains both the liquid’s vapor and the trapped air.
2. Record the gas volume and temperature as the system is brought through the experimental temperature range.
3. Use the lowest-temperature condition, near the liquid’s melting or freezing point, as an approximation where the liquid’s vapor pressure is nearly zero. Treat the measured gas pressure at this condition as pressure due to the trapped air alone.
4. Calculate the moles of trapped air at that condition using the ideal gas law, (PV=nRT), taking its pressure to be atmospheric pressure.
5. Calculate the air pressure at each experimental temperature using the same fixed amount of air and the measured temperature and volume:

[ P_{\text{air}}=\frac{nRT}{V}. ]

1. Find the liquid’s vapor pressure at each temperature using the pressure balance:

   • Total pressure is atmospheric pressure.
   • Total pressure equals the air pressure plus the liquid’s vapor pressure.
   • Therefore,

   [ P_{\text{vap}}=P_{\text{atm}}-P_{\text{air}}. ]

2. Plot (\ln(P_{\text{vap}})) against (1/T). The slope gives the enthalpy of vaporization through the Clausius–Clapeyron relationship.

The instructor notes that the same general method can be used with a suitable solvent other than water, provided its vapor pressure can be measured over the chosen temperature range.

Speakers and Sources

• Instructor/lecturer: Explains vapor pressure, the Clausius–Clapeyron relationship, and the proposed laboratory calculations. No name is given.
• Students in the CHEM 1811 class: Addressed by the instructor; no individual student speakers are identified.