Properties of gases
The kinetic theory of gases is a topic that can explain many everyday observations. Have you ever
wondered why water boils faster at higher altitudes? Or why inflatable pool toys seem flat after sitting in a
cold garage? How about why you can smell a candle all throughout the house? All of these phenomena
and many more can be explained by the kinetic theory of gases.
The kinetic theory of gases (also known as kinetic-molecular theory) explains the behavior of a
hypothetical ideal gas. According to this theory, gases are made up of tiny particles in random, straight line
motion. They move rapidly and continuously and make collisions with each other and the walls. This was
the first theory to describe gas pressure in terms of collisions with the walls of the container, rather than
from static forces that push the molecules apart. Kinetic theory also explains how the different sizes of the
particles in a gas can give them different, individual speeds.
All gases share common physical properties. Like liquids, gases freely flow to fill the container they are in.
But while liquids have a defined volume, gases have neither a defined volume nor shape.
Unlike liquids and solids, gases are highly compressible. These common properties relate to a unique
characteristic of gases: considering the interatomic distance, Gas molecules are incredibly far apart and
rarely interact with each other. In solids, the attractive and repulsive forces between molecules—the
intermolecular forces—are so strong they lock the solid into a fixed shape and size. In liquids, the
intermolecular forces are weaker, and liquid molecules can move around each other. But a liquid's
molecules are still close enough that intermolecular forces affect nearby molecules. A gas’s molecules are
so far apart that the intermolecular forces are negligible. Because gas molecules don’t interact with one
another, gases don’t exist as different types like liquids and solids do. The different types of liquids and
solids (such as molecular and network solids) have properties that reflect the unique ways their molecules
interact. As a result, all gases share some common behaviors. We can understand how any gas—whether
it’s helium or carbon monoxide—behaves by understanding the laws governing gas behavior.
Class work: Explain extensively why all gases share some common behaviors at STP.
Postulates
Kinetic theory makes many assumptions in order to explain the reasons gases act the way they do.
According to kinetic theory:
1. Gases consist of particles in constant, random motion. They continue in a straight line until they
collide with something—usually each other or the walls of their container.
2. Particles are point masses with no volume. The particles are so small compared to the space
between them, that we do not consider their size in ideal gases.
3. No molecular forces are at work. This means that there is no attraction or repulsion between the
particles.
4. Gas pressure is due to the molecules colliding with the walls of the container. All of these collisions
are perfectly elastic, meaning that there is no change in energy of either the particles or the wall
upon collision. No energy is lost or gained from collisions.
5. The time it takes to collide is negligible compared with the time between collisions.
6. The kinetic energy of a gas is a measure of its Kelvin temperature. Individual gas molecules have
different speeds, but the temperature and kinetic energy of the gas refer to the average of these
speeds.
7. The average kinetic energy of a gas particle is directly proportional to the temperature. An
increase in temperature increases the speed in which the gas molecules move.
8. All gases at a given temperature have the same average kinetic energy.
, 9. Lighter gas molecules move faster than heavier molecules.
The gas laws
Over the past four centuries, scientists have performed many experiments to understand the common
behaviors of gases. They have observed that a gas's physical condition—its state—depends on four
variables: pressure (P), volume (V), temperature (T), and amount (n, in moles; The relationships between
these variables are now known as the gas laws, which describe our current knowledge about how gases
behave on a macroscopic level.
But the relationships behind the gas laws weren't obvious at first—they were uncovered by many
scientists examining and testing their ideas about gases over many years.
Gas pressure: We now understand that air is a gas made of physical molecules . As these molecules move
about inside a container, they exert force—known as pressure—on the container when they ricochet off
its walls. Thanks to this behavior, we can inflate car tires, rubber rafts, and Macy’s Day Parade balloons
with gases. However, the idea that air is a substance made of molecules that exert pressure would have
been a strange idea to scientists before the 17th century. Along with fire, water, and earth, air was
generally considered a fundamental substance, and not one made up of other things.
Boyle's law: Boyle’s data showed that when air was squeezed to half its original volume, it doubled its
pressure. In 1661, Boyle published his conclusion that air’s volume was inversely related to its pressure.
This observation about air’s behavior—and therefore, gas behavior—is a critical part of what we now call
Boyle’s law.
Boyle’s law states that so long as temperature is kept constant, the volume (V) of a fixed amount of gas is
inversely proportional to its pressure (P)
Equation 1a
V ∝ 1/P
Boyle's law states that so long as temperature is kept constant, the volume of a fixed amount of gas is
inversely proportional to the pressure placed on the gas.
Boyle’s law can also be written as:
Equation 1b
V × P = constant
For a fixed amount of gas at a fixed temperature, this constant will be the same, even if the gas’s pressure
and volume change from (P1, V1) to (P2, V2), because volume decreases as pressure increases.
Therefore, P1 x V1 must equal the constant, and P2 x V2 must also equal the constant. Because they both
equal the same constant, the gas’s pressure and volume under two different conditions are related like
this:
Equation 2
, P1×V1=P2×V2
Considering helium balloon, Boyle’s law means that if you took the balloon deep under the ocean, it would
shrivel because the pressure is very high and the helium would significantly decrease in volume. And if you
took the balloon to the top of Mount Everest, the ballon would get even bigger (and might even pop!)
because the atmospheric pressure is low and the helium would increase in volume.
Charles's law: states that when pressure is kept constant, a fixed amount of gas linearly increases its
volume as its temperature increases
Equation 3a
V∝T
Charles's Law states that when pressure is kept constant, a fixed amount of gas linearly increases its
volume as its temperature increases.
Charles’s law can also be understood as:
Equation 3b
V T ∝ constant
For a fixed amount of gas at a fixed pressure, this constant will be the same, even if the gas’s volume and
temperature change from (V1, T1) to (V2, T2). Therefore, V1/T1 must equal the constant, and V2/T2 must also
equal the constant. As a result, the gas’s temperature and volume under different conditions are related
like this:
Equation 4
V1T1=V2T2
This means that if we took the Snoopy balloon to the North Pole, the balloon would shrink as the helium
cooled and decreased in volume. However, if we took the balloon to a hot tropical island and the helium’s
temperature increased, the helium would increase in volume, expanding the balloon.
Avogadro's law: Avogadro’s law is based off of Avogadro’s hypothesis. Avogadro’s law states that at
a constant pressure and temperature, a gas’s volume (V) is directly proportional to the number
of molecules (n, in moles) (Figure 6):
Equation 5
V∝n
The kinetic theory of gases is a topic that can explain many everyday observations. Have you ever
wondered why water boils faster at higher altitudes? Or why inflatable pool toys seem flat after sitting in a
cold garage? How about why you can smell a candle all throughout the house? All of these phenomena
and many more can be explained by the kinetic theory of gases.
The kinetic theory of gases (also known as kinetic-molecular theory) explains the behavior of a
hypothetical ideal gas. According to this theory, gases are made up of tiny particles in random, straight line
motion. They move rapidly and continuously and make collisions with each other and the walls. This was
the first theory to describe gas pressure in terms of collisions with the walls of the container, rather than
from static forces that push the molecules apart. Kinetic theory also explains how the different sizes of the
particles in a gas can give them different, individual speeds.
All gases share common physical properties. Like liquids, gases freely flow to fill the container they are in.
But while liquids have a defined volume, gases have neither a defined volume nor shape.
Unlike liquids and solids, gases are highly compressible. These common properties relate to a unique
characteristic of gases: considering the interatomic distance, Gas molecules are incredibly far apart and
rarely interact with each other. In solids, the attractive and repulsive forces between molecules—the
intermolecular forces—are so strong they lock the solid into a fixed shape and size. In liquids, the
intermolecular forces are weaker, and liquid molecules can move around each other. But a liquid's
molecules are still close enough that intermolecular forces affect nearby molecules. A gas’s molecules are
so far apart that the intermolecular forces are negligible. Because gas molecules don’t interact with one
another, gases don’t exist as different types like liquids and solids do. The different types of liquids and
solids (such as molecular and network solids) have properties that reflect the unique ways their molecules
interact. As a result, all gases share some common behaviors. We can understand how any gas—whether
it’s helium or carbon monoxide—behaves by understanding the laws governing gas behavior.
Class work: Explain extensively why all gases share some common behaviors at STP.
Postulates
Kinetic theory makes many assumptions in order to explain the reasons gases act the way they do.
According to kinetic theory:
1. Gases consist of particles in constant, random motion. They continue in a straight line until they
collide with something—usually each other or the walls of their container.
2. Particles are point masses with no volume. The particles are so small compared to the space
between them, that we do not consider their size in ideal gases.
3. No molecular forces are at work. This means that there is no attraction or repulsion between the
particles.
4. Gas pressure is due to the molecules colliding with the walls of the container. All of these collisions
are perfectly elastic, meaning that there is no change in energy of either the particles or the wall
upon collision. No energy is lost or gained from collisions.
5. The time it takes to collide is negligible compared with the time between collisions.
6. The kinetic energy of a gas is a measure of its Kelvin temperature. Individual gas molecules have
different speeds, but the temperature and kinetic energy of the gas refer to the average of these
speeds.
7. The average kinetic energy of a gas particle is directly proportional to the temperature. An
increase in temperature increases the speed in which the gas molecules move.
8. All gases at a given temperature have the same average kinetic energy.
, 9. Lighter gas molecules move faster than heavier molecules.
The gas laws
Over the past four centuries, scientists have performed many experiments to understand the common
behaviors of gases. They have observed that a gas's physical condition—its state—depends on four
variables: pressure (P), volume (V), temperature (T), and amount (n, in moles; The relationships between
these variables are now known as the gas laws, which describe our current knowledge about how gases
behave on a macroscopic level.
But the relationships behind the gas laws weren't obvious at first—they were uncovered by many
scientists examining and testing their ideas about gases over many years.
Gas pressure: We now understand that air is a gas made of physical molecules . As these molecules move
about inside a container, they exert force—known as pressure—on the container when they ricochet off
its walls. Thanks to this behavior, we can inflate car tires, rubber rafts, and Macy’s Day Parade balloons
with gases. However, the idea that air is a substance made of molecules that exert pressure would have
been a strange idea to scientists before the 17th century. Along with fire, water, and earth, air was
generally considered a fundamental substance, and not one made up of other things.
Boyle's law: Boyle’s data showed that when air was squeezed to half its original volume, it doubled its
pressure. In 1661, Boyle published his conclusion that air’s volume was inversely related to its pressure.
This observation about air’s behavior—and therefore, gas behavior—is a critical part of what we now call
Boyle’s law.
Boyle’s law states that so long as temperature is kept constant, the volume (V) of a fixed amount of gas is
inversely proportional to its pressure (P)
Equation 1a
V ∝ 1/P
Boyle's law states that so long as temperature is kept constant, the volume of a fixed amount of gas is
inversely proportional to the pressure placed on the gas.
Boyle’s law can also be written as:
Equation 1b
V × P = constant
For a fixed amount of gas at a fixed temperature, this constant will be the same, even if the gas’s pressure
and volume change from (P1, V1) to (P2, V2), because volume decreases as pressure increases.
Therefore, P1 x V1 must equal the constant, and P2 x V2 must also equal the constant. Because they both
equal the same constant, the gas’s pressure and volume under two different conditions are related like
this:
Equation 2
, P1×V1=P2×V2
Considering helium balloon, Boyle’s law means that if you took the balloon deep under the ocean, it would
shrivel because the pressure is very high and the helium would significantly decrease in volume. And if you
took the balloon to the top of Mount Everest, the ballon would get even bigger (and might even pop!)
because the atmospheric pressure is low and the helium would increase in volume.
Charles's law: states that when pressure is kept constant, a fixed amount of gas linearly increases its
volume as its temperature increases
Equation 3a
V∝T
Charles's Law states that when pressure is kept constant, a fixed amount of gas linearly increases its
volume as its temperature increases.
Charles’s law can also be understood as:
Equation 3b
V T ∝ constant
For a fixed amount of gas at a fixed pressure, this constant will be the same, even if the gas’s volume and
temperature change from (V1, T1) to (V2, T2). Therefore, V1/T1 must equal the constant, and V2/T2 must also
equal the constant. As a result, the gas’s temperature and volume under different conditions are related
like this:
Equation 4
V1T1=V2T2
This means that if we took the Snoopy balloon to the North Pole, the balloon would shrink as the helium
cooled and decreased in volume. However, if we took the balloon to a hot tropical island and the helium’s
temperature increased, the helium would increase in volume, expanding the balloon.
Avogadro's law: Avogadro’s law is based off of Avogadro’s hypothesis. Avogadro’s law states that at
a constant pressure and temperature, a gas’s volume (V) is directly proportional to the number
of molecules (n, in moles) (Figure 6):
Equation 5
V∝n
