Molecular Motion

Kinetic-Molecular Theory

The ideal gas equation

pV = nRT

Has been presented as a compliation of empirical observation, i.e. the historically significant Gas Laws, but does The Ideal Gas equation have some deeper, more fundamental meaning?

The Kinetic-Molecular Theory ("the theory of moving molecules"; Rudolf Clausius, 1857)

  1. Gases consist of large numbers of molecules (or atoms, in the case of the noble gases) that are in continuous, random motion.  Usually there is a great distance between each other, so the molecules travel in straight lines between abrupt collisions at the walls and between each other.  These collisions randomize the motion of the molecules. Most of the collisions between molecules are binary, in that only two molecules are involved.
  2. The volume of the molecules of the gas is negligible compared to the total volume in which the gas is contained.  A common bond length between atoms is about 10-10 m or 1 Angstrom. Small molecules are therefore on the order of 10 Angstroms in diameter, or less than 10-24 Liters in Molecular Volume, quite tiny indeed!  Remember, however that there can be a great many molecules in the sample of gas, perhaps on the order of a mole, or 6 x 1023.  So that when concentrations of molecules exceed about 1 mol/liter, then the approximation that the volume of ALL the molecules in the container is much less than the volume of the container itself, fails.  In the case of an ideal gas, we will assume that molecules are point masses, i.e., the volume of a mole of gas molecules (as if they were at rest) is zero, so molecular and container volumes never become comparable.
  3. Attractive forces between gas molecules are negligibleWe know that if these forces were significant, the molecules would stick together.  This happens when it rains and gaseous water molecules stick together to form a liquid.  Water vapor is a condensible gas, and this shows us that gas molecules are sticky, but at a high enough temperature they form only a permanent gas, because their stickiness can be considered negligible.  We will assume that in an ideal gas, molecular attractive forces are not just small, but identically zero. 
Consequences: The Molecular picture of Pressure

Absolute Temperature

Molecular Speed, Averaging, and The Maxwell-Boltzmann Speed Distribution

The following graph shows the Distribution of Speeds for Gases, i.e.the fraction of the sample of gas molecules that have a given speed is shown by the height of the curve above the speed axis.  There are no molecules exactly at rest.

Where does this number lie on our Graph of the speed distribution?  To find out, the root mean square 'average' of the distribution must be taken (defined) in a specific way.  It is defined exactly how it sounds.  First your square all the speeds, then average those numbers, and take the square root of that average.  The mean of a distribution is just the average of all the numbers in the distribtion.  The most probable value of a distribution is also exactly what it sounds like, the speed of that the largest fraction of molecules are travelling.  In general, the mean, the root mean square and the most probable value in a distribution are all different.

A Note on Distributions a simple numerical example:

Suppose we have four molecules in our gas sample. Their speeds are 3.0, 4.5, 5.2 and 8.3 m/s.

The rms speed as well as the entire distribution of speeds of gas molecules are a function of temperature. Below, the blue line is a cold gas and the red line is a hot gas. Note that the rms speed, u as well as the entire speed distribution changes with temperature for a given gas.
The rms speed for a given speed distribution (which is determined by the temperature and molecular weight of the gas) is greater in magnitude than the most probable speed or the mean speed.

Trick question: What is the mean velocity of the molecules in a gas at any temperature?

Gas Laws and Kinetic Theory

The Ideal Gas Equation of State follows directly from the Kinetic Theory of Gases. Here is a Pseudo-Derivation

Molecular Effusion and Diffusion

Kinetic-molecular theory states that the average kinetic energy of a mole of molecules molecules is proportional to absolute temperature, and the proportionality constant is R, the universal gas constant

(1/2)Mu2 = (3/2)R T   =   Molar Kinetic (translational) Energy of the gas

where M is the molar mass in kg/mole, R is the gas constant in J/K.mole, and T is the absolute temperature in K.

Numerical Example:

Calculate the rms speed, u, of an N2 molecule at room temperature (25°C) Be careful of your UNITS!

T = (25+273)K = 298K
M = 28 g/mol = 0.028 kg/mol
R = 8.314 J/mol K = 8.314 kg m2/s2 mol K
u = 515 m/s

Note: this is equal to 1,150 miles/hour!


The escape of a gas through a tiny pore or pinhole in its container is called EFFUSION.

The effusion rate, r, has been found to be inversely proportional to the square root of its molar mass: (Why?)

Thus, comparison of the effusion rates of two gases with different masses will follow the relation:

This effect was observed in the 19th century by Graham and is sometimes called GRAHAM's LAW

A note on Rates and Times
The effusion time (the time it takes for a certain amount of gas to escape a vessel) is inversely proportional to the effusion rate (the amount of gas effusing from the hole per unit time). Be careful that you understand whether it is a rate or a time that you are calculating.

Gas may effuse, but for this to happen a molecule must pass through a pore or pinhole and escape to the outside. In effect, a molecule must 'collide' with an escape hole. The number of such collisions will be linearly proportional to the average speed of the molecules in the gas and thus the effusion rate. The effusion time should be inversely proportional to the average speed of the molecules or proportional to the square root of the ratio of the molecular masses..

The ratio of effusion rates, ri, for two gases labelled by i, is proportional to the ratio of the RMS speeds of the gases, ui


Similarly to effusion, the process of diffusion is the spontaneous intermingling (mixing) of dissimilar gases (fluids) that are initially spatially separated. If you put a drop of ink in a glass of water and you see the ink gradually spread out to fill the glass, this is diffusion
The average distance traveled by a molecule between collisions with another molecule is called the mean free path
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