Chapter 1
Introduction to Gases

Gases represent a state of matter that has no fixed shape or fixed volume, which consist of tiny, energetic particles, i.e., atoms or molecules, that are widely spaced (Figure 1.1). Compared to the other states of matter, solids and liquids, gases have a much lower density, i.e., they have a small mass per unit volume, because there is a great deal of empty space between gas particles. At room temperature and pressure, the gas inside a container occupies only 0.1% of the total container volume. The other 99.9% of the total volume is empty space (whereas in liquids and solids, about 70% of the volume is occupied by particles). Gas particles move very fast and collide with one another, causing them to diffuse, or spread out, until they are evenly distributed throughout the volume of their container. You will never see only half of a balloon filled with air.

Figure 1.1 Gas atoms or molecules are constantly moving and colliding with one another.

Although both liquids and gases take the shape of their containers, gases differ from liquids in that there is so much space between gas molecules that they offer little resistance to motion and can be compressed to smaller and smaller volumes. As seen in Figure 1.2, as a gas is compressed, the molecules making up the gas get closer together and create a higher internal pressure.

Figure 1.2 As gas is compressed, the gas molecules get closer together.

Hydrogen is the lightest known gas. Any balloon filled with hydrogen gas will float in air if the total mass of its container is not too great. Helium gas is also lighter than air and has 92% of the lifting power of hydrogen. Today all airships, i.e., blimps, use helium instead of hydrogen because it offers almost the same lifting power and is not flammable.

Gases can be monatomic, diatomic, and polyatomic. Monatomic gases are gases composed of single atoms, diatomic gases are those composed of two atom molecules, and polyatomic gases are those made up of molecules with more than two atoms. Noble gases such as helium, neon, argon, etc., are normally found as single atoms, since they are chemically inert. Gases such as nitrogen (N2), oxygen (O2), and carbon monoxide (CO) tend to be found as diatomic molecules (Figure 1.3). Carbon dioxide (CO2), and methane (CH4) are examples of polyatomic gas molecules (Figure 1.3).

Figure 1.3 Oxygen, nitrogen, and carbon monoxide are examples of diatomic molecules. Carbon dioxide, water, nitrogen monoxide, methane, sulfur dioxide, and ozone are examples of polyatomic molecules.

Gases can be found all around us. In fact, the earth's atmosphere is a blanket of gases composed of nitrogen (78%), oxygen (21%), argon (1%), and then trace amounts of carbon dioxide, neon, helium, methane, krypton, hydrogen, nitrous oxide, xenon, ozone, iodine, carbon monoxide, and ammonia.

Because of the large distances between gas particles, the attractions or repulsions among them are weak. The particles in a gas are in rapid and continuous motion. For example, the average velocity of nitrogen molecules, N2, at 68 °F is about 1640 ft/s. As the temperature of a gas increases, the particles' velocity increases. The average velocity of nitrogen molecules at 212 °F is about 1886 ft/s. The particles in a gas are constantly colliding with the walls of the container and with each other. Because of these collisions, the gas particles are constantly changing their direction of motion and their velocity. In a typical situation, a gas particle moves a very short distance between collisions. For example, oxygen, O2, molecules at normal temperatures and pressures move an average of 0.000003937 inches between collisions.

1.1 Ideal Gases

Scientists often simplify the model of gases by imagining the behavior of an ideal gas. An ideal gas differs from a real gas in that the particles are assumed to be point masses, that is, particles that have a mass but occupy no volume. It is also assumed that there are no attractive or repulsive forces at all between the particles. When all these assumptions are incorporated into a gas model, the "ideal gas model" is obtained. As the name implies, the ideal gas model describes an "ideal" of gas behavior that is only approximated by reality. Nevertheless, the model has been proven to reasonably explain and predict the behavior of typical gases under typical conditions.

Note: Under ordinary conditions, the properties of gases predicted by the ideal gas law are within 5% of their actual values.

1.2 Properties of Gases

The ideal gas model is used to predict changes in four related gas properties: volume, number of particles, temperature, and pressure. Volumes of gases are usually described in cubic feet, ft3, or cubic meters, m3, and numbers of particles are usually described in moles.

1.3 Temperature

Temperature is a physical quantity expressing how hot or cold a system of atoms or physical object is. Technically, temperature is the proportional measure of the average kinetic energy related to the random motions of the constituent particles of matter in a system. Temperature is an important property of a system because it is an indication of the direction in which heat energy will spontaneously flow. Remember that heat energy always flows from a hotter body (one at a higher temperature) to a colder body (one at a lower temperature).

Temperature is a measure of the total heat energy in a system.

Gas temperatures can be measured with thermometers, infrared guns, and thermocouples. Readings can be reported in degrees Fahrenheit, °F, or Celsius, °C. However, engineers generally use Rankine, or Kelvin temperatures for calculations.

1.4 Pressure

Remember that gases have no definite shape or volume; they tend to fill whatever container they are in. They can compress and expand and have extremely low densities when compared to a liquid or solid. Combinations of gases tend to mix together spontaneously; that is, they form gas mixtures. Air, for example, is a solution of mostly nitrogen and oxygen. Any understanding of the properties of gases must be able to explain the properties of gas mixture.

The kinetic theory of gases indicates that gas particles are always in motion and are colliding with other particles and the walls of the container holding them. Although collisions with container walls are elastic (i.e., there is no net energy gain or loss because of the collision), a gas particle does exert a force on the wall during the collision. Each time a gas particle collides with and ricochets off one of the walls of its container, it exerts a tiny force against the wall. The accumulation of all these forces distributed over the area of the walls of the container causes something we call pressure. Pressure (P) is defined as the force of all the gas particle-wall collisions divided by the area of the wall:

In English units, pressure is measured in psi, or pounds per square in. The formal, SI-approved unit of pressure is the pascal (Pa), which is defined as 1 N/m2 (one newton of force over an area of one square meter). However, this is usually too small in magnitude to be useful. A common unit of pressure is the atmosphere (atm), which was originally defined as the average atmospheric pressure at sea level.

1.5 Gas Laws

When seventeenth-century scientists began studying the physical properties of gases, they noticed simple relationships between some of the measurable properties of gases. For example, scientists noted that for a given quantity of gas, usually expressed in units of moles, i.e., number of molecules [n] in a system, if the temperature (T) of the gas is kept constant, pressure and volume are related: as one variable increases, the other variable decreases. Conversely, as one variable decreases, the other variable increases. Therefore, we say that pressure and volume are inversely related.

take pressure (P) and volume (V), for example:

There is more to it, however: pressure and volume of a given amount of gas at a constant temperature are numerically related. If you take the pressure value and multiply it by the volume value, the product is a constant for a given amount of gas at a constant temperature:

(1.1)

If either volume or pressure changes while the amount and temperature stays the same, then the other property must change so that the product of the two properties still equals that same constant. That is, if the original conditions are labelled P1 and V1 and the new conditions are labelled P2 and V2, we have

(1.2)

where the properties are assumed to be multiplied together. Leaving out the middle part, we have simply:

(1.3)

This equation is an example of a gas law. A gas law is a simple mathematical formula that allows you to model, or predict, the behavior of a gas. This particular gas law is called Boyle's Law, after the English scientist Robert Boyle, who first announced it in 1662. Figure 1.4 shows two representations of what Boyle's Law describes.

Figure 1.4 Starting with a piston having a given pressure and volume (far right piston), the volume continuously decreases as the applied pressure increases. If you plot pressure (P) as a function of the volume (V) for a given amount of gas at a certain temperature, you will get a plot that...