Study Guide

Solids, Liquids, and Gases - Gases

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The most common terms associated with gases are temperature and pressure. Temperature is a measurement of the average kinetic energy of the molecules in a system, usually measured with a thermometer and expressed in Kelvin (K). While working through exercises, if you encounter a problem that has temperature in Celsius, convert it to Kelvin in the following manner:

Kelvin = Celsius + 273.15

A practice exercise might also use the term standard temperature. This is defined as 0 oC or 273 K. Notice the units for Kelvin are not expressed in degrees. In fact, the General Conference on Weights and Measurements abolished the use of degrees when referring to Kelvin in 1967. If you think chemistry is tedious, just be glad you don't have to sit through that conference. Make sure you omit the degree symbol when using K.

Pressure is defined as the force per unit area, usually measured with a barometer. There are three common units for pressure: atmospheres (atm), millimeters of mercury (mm Hg), and Pascals (Pa). Standard pressure is defined as 1 atm of 760.0 mm Hg. A common practice in chemistry is to refer to conditions as STP, or standard temperature and pressure. When we encounter this, it means the reaction was run at 273 K and 1 atm.

The gaseous state is the most simple and least fixed phase of matter. It has no definite volume or shape. Because of this, gases are subject to pressure, volume, and temperature changes, all of which affect the overall properties of the gas. The particles in a gas move so rapidly and are so loosely arranged that they will fill any shape in which they are put. Thank goodness, because where would we be without balloon animals?

Need a break? Make a monkey.

With its territory vast and its occupants few, a gas consists of mainly empty space. The gas particles, therefore, are pretty much on their own. This is the basis of the ideal gas model. The ideal gas in an approximation is a theoretical description of a gas state and is not a description of any specific substance. There are no ideal gases in real life. Scientists constructed the ideal gas model to simplify the gas phase and make calculations more manageable.

An ideal gas is described by the following characteristics, known collectively as the kinetic molecular theory (KMT):

1. Contains tiny, discrete particles that have mass but virtually no volume

2. The particles are in constant, rapid, and random motion

3. No attractive forces exist between the particles

4. No attractive forces exist between the particles and their container

5. When the particles collide, energy is conserved

6. No energy is lost when a particle collides with the container

These characteristics are the foundation of the ideal gas and the gas laws that describe gas behavior, namely: Boyle's Law, Charles' Law, Avogadro's Law, the Combined Gas Law, the Ideal Gas Law, Dalton's Law of Partial Pressures, and finally Graham's Law.

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It's the gas man.

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