GAS PRESSURE

Unlike most liquids, gases can be compressed. This is why it is possible to fill up hundreds of balloons with a single, small tank of helium gas and why it is possible for a scuba diver to breathe for a long time from a single small tank of air. Imagine a container whose volume (in meters cubed) is equal to V.

Suppose that there are N moles of atoms of a particular gas inside this container, which is surrounded by a perfect vacuum. We can say certain things about the pressure P, in newtons per meter squared, that the gas exerts outward on the walls of the container. First, P is proportional to N, provided that V is held constant. Second, if V increases while N remains constant, P will decrease. These things are apparent intuitively.
There is another important factor—temperature—involved when it comes to gases under pressure when they expand and contract. The involvement of temperature T, generally measured in degrees above absolute zero (the absence of all heat), is significant and inevitable in gases. When a parcel of gas is compressed, it heats up; when it is decompressed, it cools off. Heating up a parcel of gas will increase the pressure, if all other factors are held constant, and cooling it off will reduce the pressure. The behavior of matter, especially liquids and gases, under conditions of varying temperature and pressure is a little complicated, so the entire next chapter is devoted to this subject.