|Lecture Notes: 29 October||
Example: A student ignores the warning labels and throws an empty (no liquid left, no spray) can of hair spray into his campfire. Assuming an ambient temperature of 25 °C and atmospheric pressure of 7.20 x 102mmHg, and a temperature in the coals of 600 °C, find the pressure in the can in the fire, assuming it doesn't burst or expand.
Example: One of the student's colleagues on this ill fated trip tossed an "empty" 0.500 L propane cylinder into the fire. Unfortunately, 3.50 g of propane remained in the cylinder. What pressure would be reached in the cylinder assuming no deformation and no bursting at 550 °C (assume 3 sig figs for the temp.).
Now we can expand our view of gas behavior with two additional observations, both consistant with our picture of gas behavior.
Equal volumes of different gases at the same temperature (T) and pressure (P) have the same number of particles (same number of moles).
Example: 2.40 L of ethene gas (C2H4) is combined with 7.35 L of oxygen and ignited. If all volumes of reactants and products are measured at the same temperature and pressure (above 100 °C - so water is a vapor), calculate the volume of each substance after the reaction is complete.
The pressure of a gas is independent of the presence of other gases.
One of the most common applications of Dalton's Law of Partial Pressures is the determination of gas pressures above water. We will see in lab that the vapor pressure (gas pressure) of water is fixed by the temperature, and that there will always be a contribution to the total pressure by water when it is present.
Example: 50.0mL of oxygen is collected over water from a specimen of Anacris water weed illuminated by controlled lighting. If the temperature is 20.0°C and the pressure of the collected gas is 760.5 mmHg how many moles of oxygen were collected? (the vapor pressure of water = 17.5 mmHg at 20.0°C).
We have been looking at the various properties of gases, now we want to look at a theory to explain those behaviors. A simple model is the kinetic-molecular theory. There are four basic postulates:
Note that for kinetic energy, KE = 1/2 mv2, so v (velocity) varies as the square root of the mass (m1/2). Notice also that the energy increases with the square of the velocity. (This is why an accident at 60 mph is much worse that one at 30 mph - four times as much energy is involved!).
© R A Paselk
Last modified 29 October 2009