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Try the temperature control (slider) to see what changes in the gas

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Gases, like other states of matter (solids , liquids and plasma), consists of molecules or atoms . In gases, the molecules and/or atoms move freely in all directions and bounce of the walls of the container, as it is shown on the animation. Because every atom (or molecule) in the gas can move freely, it means that a gas takes up all the space available to it, so that it fills the container (unlike a liquid, or solid, that rests at the bottom of the container).

An increase in temperature correspond to an increase in the speed of the atoms (or molecules) , as seen on the animation; this will also cause more frequent collisions with the walls and consequently, an increase in pressure. If you increase the number of atoms the pressure will also increase .

Pressure, which is a macroscopic effect, is caused by atoms (or molecules) hitting the walls of the container (a microscopic effect). Obviously, an increase on the number of particles will also cause an increase in pressure.

If the size of the container increases, the pressure of the gas diminishes.

Temperature is also a macroscopic measurement. Isolated atoms or molecules don't have a temperature.

Gases may be atomic or molecular .A common example of an atomic gas is Helium; all inert gases (located at the last column of the periodic table on the right) are atomic gases. In this case each particle of the gas is only an atom (isolated).

Gases like Oxygen or Nitrogen (which form the terrestrial atmosphere) are molecular. Two oxygen atoms form the oxygen molecule and two nitrogen atoms form the nitrogen molecule (gas molecules may also be composed of more than two atoms).

Molecular gases exist when the intermolecular forces are weak.

Ideal gas equation:

PV = n RT

P : pressure
V : volume
R : constant
T : temperature
n : number of moles

If the number of particles is kept constant, that ideal gas equation simplifies to:

PV = T

The behaviour of each particle on the gas can be modeled using Newton's equations, so that the ideal gas equation shown above can be deduced by first principles (that is called kinetic theory of gases).


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In animation 1, above, we see a gas inside a container which is smaller than the one in animation 2 (below).

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As a result, the atoms collide much more often with the walls of the container; this fact is perceived macroscopically as an increase in pressure (the atoms moving and colliding constitute the macroscopic view).

The animations just show 3 atoms in the container but in a normal situation you will have many billions of them and they will be very small .