http://www.ck12.org Chapter 14. Thermodynamics
Introduction
In the chapter covering temperature and heat, we discussed the kinetic theory of temperature for an ideal gas. The
kinetic theory stated that the temperature of such a gas is directly related to the kinetic energy of the molecules of the
gas, which, in turn, is equal to the internal energy of the gas. The more we heat a gas, the faster the random motion
of its molecules and the greater the increase in temperature and pressure. This seems perfectly reasonable if the gas
is confined so that its volume cannot increase. Heat an empty glass bottle with a cork at the top for a sufficiently
long time and the molecules of the air in the bottle gain enough kinetic energy to eventually “pop” the cork out of
the bottle. In other words, the pressure (force per unit area) becomes great enough to force the cork out of the bottle.
But what happens if we heat a gas and allow the volume to increase, or compress a gas while cooling it? In this
chapter, we explore the relationships between the mass, volume, pressure, and temperature of ideal gases which are
confined, compressed, and expanded.
Atomic Mass Units
Before we continue, it will be helpful to discuss how the masses of different atoms (and molecules) are compared to
each other.
The idea that all matter is composed of tiny building blocks called atoms goes back to antiquity. It was only during
the last two hundred years or so, however, that some measurements could be made suggesting the truth of the idea.
Measurements made when different substances combined led to the conclusion that the ratios of atomic masses
tended to be simple proportions and that hydrogen had the smallest mass of all atoms. Compared to a hydrogen
atom, for example, a carbon atom had 12 times the mass, an oxygen atom-16, and nitrogen-14.
Today we define the mass of one-twelfth of a carbon atom as one unified atomic mass unit(amu)or(u)
1 u= 1. 66 × 10 −^27 kg
Using this definition, a carbon atom has a mass of 12.0000 u, and a hydrogen atom has a mass of 1.0078 u.
The Mole
A basic unit in SI system is themole(mol). The mole is defined as the amount of a substance that contains the same
number of atoms or molecules as exactly 12 grams of carbon-12. For example, it is found that 16 grams of oxygen
contains the same number of atoms as does 12 grams of carbon-12. Note that carbon-12 is the most common isotope
of carbon. We’ll have more to say about isotopes later.
Molar Mass
Molar mass is the mass of one mole of a substance.
Using the atomic mass of a compound, we can quickly determine the mass of one mole of the compound. For
example, carbon monoxide (CO) is a chemical compound composed of one carbon atom and one oxygen atom. The
unified atomic mass of carbon is 12 and the unified atomic mass of oxygen is 16. The total unified atomic mass
number for carbon monoxide is therefore:
12 + 16 = 28 u
The molar mass of carbon dioxide is numerically equal to the atomic mass number, but in grams. Therefore, carbon
dioxide has a molar mass of 28 grams.
Another definition of the mole can be stated as: The number of grams of a substance that is equal to its molar mass.