In Chapter 1 and 2 we will discuss in detail the three states of matter. Here is a macroscopic view of each state and a microscopic view.

If we look at a periodic table which identifies the phase of each element at standard conditions we see that the dominate phase is solid followed by the gas phase and finally the liquid phase.

The elements which are gases include, hydrogen, helium, oxygen, nitrogen, fluorine, chlorine and all the noble gases. I'll expect you to know the standard state phase of every element in the periodic table. Now you might be thinking, holy cow, are you nuts, but remember 2 of the elements are liquids and 10 of the elements are gases. So it is not so bad. Also I'll expect you to know the color of each of the gaseous elements.

Under appropriate conditions substances which are normally liquids or solids can exist in the gas phase. In such cases the gaseous phase of the substance is called vapor. For example, one of the two elements which exists as a liquid is bromine.

The three phases of matter differ very obviously from one another. A gas is differentiated from other matter by two characteristics: 1) it has no definite shape and 2) it has no definite volume, but expands to occupy its container.

The liquid state has a definite volume, but no definite shape. The liquid flows and when poured from one container to another takes the shape of the container. Liquids have no definite shape.

A solid does not flow but has a definite volume and shape. To change the shape requires the application of considerable force.

Perhaps the most familiar property of a gas is its compressibility. The properties of gases such as compressibility, mixing and diffusion occur because of the large separation between individual molecules. In the solid phase individual molecules are very close together, while in the gas phase the separation between molecules is very large.

So our goal in the first part of this chapter will be to develop a quanitative understanding of how temperature, volume and pressure effect the properties of a gas.

We'll begin with a discussion of pressure.

Pressure is a measure of the force of an object, or collection of particles, on a given area. Another way of saying the same thing, is pressure is the force exerted on an object divided by the area over which the force is distributed. Mathematically, we would write this relationship in the following way;

We need to make a distinction between pressure and force. Something familiar to all of us is going down to the gas station to pump up a bicycle tire. Tires require between 40 psi (pounds per square inch) to 130 psi. We know adding more air to the tire increases the pressure. The number of psi increase. The units (psi) describe the weight exerted by the gas divided by the area over which the weight is distributed. We also understand that P is proportional to the number of moles of gas in a container.

The difference between force and pressure can be described in terms of a person walking on a frozen lake. Up right, the person weight (mass x gravity) is distributed over a small area (area of shoes). If the ice is thin it may break as a result of the person standing on it. By lying down, and lying flat on the ice the individual's weight is distributed over a greater area and the pressure decreases.

Atmospheric pressure is a term that should also be familiar to you. The local weather report includes a map of the US with indicated regions of high and low pressure. This suggests that pressure is not constant. If we read the newspaper or watch a weather report the pressure is stated and it changes each day. You also know that at high altitudes (show an airplane) the pressure is lower than on the ground.

We can measure the pressure exerted by the atmosphere by filling a hollow glass tube with mercury, (a tube longer than 76. cm) and while plugging the bottom of the tube inverting it into a pan containing mercury. (Note: the mercury barometer was first developed by Evangelista Torricelli in 1643) Such a device is called a barometer (). The atmospheric pressure is obtained by measuring the distance between the surface of the mercury in the reservoir and the top of the mercury in the tube. Well you might ask, why didn't the mercury run completely out of the tube when it was inverted into the pan? The answer is the pressure exerted by the atmosphere supports the column of mercury. Another way to describe this is the weight of mercury in a column 760 mm high is equal to the weight of the air above the surface of the mercury in the pan.

We can demonstrate the presents of atmospheric pressure by the following simple experiment (QuickTime movie, 1.6 M file)(this version is displayed in a larger frame size/2.4 M file). If a soda can is partially filled with water and the water heated to boiling, nearly all of the air is swept from inside the can. If the top of the can is quickly sealed and the water in the vapor phase rapidly cooled a large difference in pressure is obtained. This can be accomplished by inverting the can and immersing it in a container of water at room temperature. . We will discuss aspects of this experiment later in the lecture. Area of the pop can is 2(pi)r(h + r) = 2(3.14)(1.125)(4.75+1.125) = 32 in2. Pressure is so the weight of the atmosphere on the can is 14.7(32) = 470 lbs.

Boyle's Law data and plot.

Charles' Law data and plot.