This chapter was actually called Thermodynamics but in class we just played around with the Ideal Gas Law,
PV = nkT,
PV = NRT
P = Pressure, V = Volume, n = number of atoms, k is the Boltzmann’s constant, and T = Temperature. The second formula is basically the same thing but instead of the number of atoms, N = the number of moles. 1 Mole contains 6.0×1023 atoms, so its just another way of talking about the number of atoms, like a dozen eggs, means 12 eggs, 1 mole of Hydrogen means 6.0×1023 H atoms. R is the Ideal Gas Constant.
Before we dug into the ideal gas law, we discussed pressure and its definition as Force per unit Area. I gave the example of stepping on someone’s toes wearing my sneakers… not too bad since my weight is spread out over the bottom of my sneaker, but, what if I was wearing spike heels and stepped on your toe with the heel?! All the students winced in imagined pain, knowing intuitively that that would be more painful. The same force applied over a tiny area has a greater pressure than if applied over a large area. We also talked about how pressure in a fluid (liquid or gas) only depends on the density of the fluid and the height of the fluid above you. The deeper you go in the ocean, the more pressure your body will experience from all water above you. Same for us standing on dry land, but under an ocean of air. As we go up in our atmosphere there is less air above us and the pressure decreases. Animals that live in the deepest part of the ocean are under extremely high pressure but their bodies are built to withstand those pressures, but if they are brought up to the surface their bodies explode because of the high pressure inside their body and sudden low pressure outside (when at the surface). Our bodies have trouble adjusting to changes in pressure as well. Divers have to be careful not to get decompression sickness or the bends – which happens when nitrogen dissolved in their blood forms bubbles when they return to the surface too quickly. This can have terrible effects on the body, including the brain.
I did a few demonstrations in this class, including letting the students try to pull apart the miniature Magdeburg hemisphere that I had. The story with this goes that the original Magdeburg hemispheres were made to demonstrate the air pump (vacuum pump) that Otto von Guericke had created. They pumped out the air and 30 horses couldn’t pull them apart! The air pressure outside the spheres was so much greater than the low pressure inside they could not be pulled apart. These smaller ones are just pushed together like suction cups to force out the air, so the pressure difference is not huge, but its enough that only the strongest students could pull them apart.
We also crushed a few soda cans with air pressure. You take an empty soda can, put a little bit of water in the bottom, maybe a cm deep. Put it on a hotplate or your stove and heat it until you see steam escaping from the top of the can. Then grab the can with tongs and quickly turn it upside into a pot of very cold ice water. The can will almost instantly collapse as the water vapor inside the can condenses leaving very little ‘air’ inside the can so the outside pressure can easily crush the can. You can see this demo and learn a lot about the ideal gas law by watching this Crash Course video from their Chemistry series.
You can also find a lot of videos on youtube where they crush slightly bigger cans.
I also had the students put little marshmallows in large syringes which we could cap. By pulling back on the syringe you increase the volume (V) of the gas, since T is constant, P must decrease (see ideal gas law above). When the pressure decreases the marshmallow expands and when you increase the pressure by pushing the syringe in (decreasing V), the marshmallow shrinks. This works because the marshmallow is filled with little pockets of air. We put a few in the microwave and watched the volume increase as the temperature increased (P was constant)… yet another example of the ideal gas law. Here’s a youtube video demonstrating the marshmallows in the syringe.
The main experiment involved Boyle’s law, which is still the ideal gas law, but with T constant you end up with PV = constant. I used a Boyle’s Law apparatus which is just a syringe embedded in wooden blocks so its easy to stack books on the syringe to increase the pressure on the volume of gas trapped in the syringe. The students started with a volume of 30mL, put the cap on the syringe (hidden by the bottom wooden block in the photo), then continued to measure the volume of the trapped air as they placed more and more books on the syringe. They recorded the applied mass, calculated the Force and Pressure on the gas, added atmospheric pressure (the pressure exerted by the air around us) and they recorded the new volume. At the end they multiplied the calculated P (from their applied mass) by the Volume and found it was indeed constant. We ran out of time in class but they should be making a graph of V vs P to put their lab book.