The Kinetic Theory of Gases
is one of the most interesting topics in Chemistry. If you come to truly understand
the concepts in this chapter, it will change the way you look at the world around you.
One of the great things about this topic is that it explains some of the phenomena
that you encounter in your everyday life. For example, have you ever inflated a pool
float until it was firm, thrown it into a cold pool, and then wondered why the float then
seemed like it was not fully inflated? Do you know why a basketball seems flat after
it has spent the night in a cold garage? What determines the time it takes to smell
the perfume of a woman who walks past you? How do hot air balloons work? How
can a small barbeque tank hold enough propane to cook with all summer long? All of
these questions can be answered by someone who has studied this chapter.
Gases have special properties
that liquids and solids don't have. The molecules that make up the gas are free to
move about, and a gas will take up the size and shape of its container. Knowing the
volume of a gas tells you very little about the quantity of matter, because any sample of
gas will fill its container. If you have a ten-gallon tank on your barbeque, it is
always technically full! In order to have an idea of the amount of matter that a
sample of gas represents, you need to know the temperature and the pressure of the gas.
Ideal
Gases - Consider how different a gas is from a solid. In a gas, the size of
the sample has very little to do with the size of the actual atoms that make up the gas
itself. Even in relatively dense gas samples, the space in between the molecules
will be much larger than the molecules themselves. When we do math problems
involving gases, we treat the particles as point masses, or particle with mass
but no volume. Ideal gases differ from real gases in another important way.
In real gases, there will be an attraction between the particles involved. These
attractions are often minor and we ignore them when we do math problems involving
gases. It is important to remember the differences between real gases and ideal
gases. It is also interesting to note that real gases will act most like ideal gases
at low pressure and high temperature, when the gas sample is less dense.
Pressure - You
may recall that pressure is defined as a force over an area. In Chemistry, pressure
is often measured in kilopascals (kPa), millimeters of mercury (mm of Hg), or atmospheres
(atm). For convenience sake, a standard atmospheric pressure has been set at 101.3
kPa, which is also equal to 760 mm of Hg and 1.0 atm. As a student of Chemistry you
should be aware of the following constants and conversions:
Standard Atmospherice Pressure = 101.3 kPa
= 760 mm of Hg = 1.0 atm
1 kPa = 7.50 mm of Hg
Temperature - Many of the thermometers that are used in
Chemistry laboratories are marked with the Celsius scale. However, when we do math
problems involving the temperature, volume and pressure of gases, we must use the Kelvin
scale. The reason for this is the fact that it is possible to have negative numbers
on the Celsius scale, and that would cause problems when measuring the volume of a gas at
low temperatures. In order to do any gas law calculations involving temperatures,
you must first convert the temperature to Kelvin. As a reminder, the conversions
between Kelvin and Celsius are shown below.
Co + 273 = K
K - 273 = Co
For convenience, standard temperature has been set at 273 Kelvin, which is
equal to 0oC. Standard temperature and pressure is abbreviated as STP.
Conditions will vary from laboratory to laboratory and from day to day. You will
often be called upon to adjust the volume of the gas that you collected in your own lab to
STP, meaning standard conditions for pressure and temperature. Remember the
information below.
STP = 101.3 kPa and 273 K (or any
equivalent values, i.e. 1 atm and 0oC)
Kinetic Theory of
Gases Quizzes |
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Kinetic Theory of
Gases Worksheets |
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Please forward all questions, comments and criticisms to Gregory L. Curran.
© Copyright 2004 Fordham Preparatory School, All Rights Reserved.
Last Modified February 07, 2008 |
