What do we assume about the volume of the actual molecules themselves in a sample of gas, compared to the bulk volume of the gas overall? Why?
One of the assumptions made in the Kinetic Molecular Theory is that the molecules have zero volume. Neglecting the volume term (and it is small) allows to to work with a much simpler equation in the general gas law of PV = nRT. There are other assumptions, too. You can read more about it here.
http://www.chm.davidson.edu/vce/kineticmoleculartheory/basicconcepts.html
In the kinetic theory of gases, it is assumed that the volume of the actual molecules themselves is negligible compared to the overall bulk volume of the gas. This assumption is made because gas molecules are considered to be point particles with no volume.
To understand why we make this assumption, let's consider the nature of gases. Gases consist of a large number of particles (atoms or molecules) that are in constant random motion. The volume of a gas is determined by the combined free space between particles.
In most cases, the sizes of gas molecules are extremely small compared to the intermolecular distances. The distance between gas molecules is generally much larger than the size of the molecules themselves. Hence, the volume occupied by the actual molecules is relatively insignificant compared to the overall volume of the gas sample.
To visualize this, imagine a container filled with gas molecules. Even though the molecules themselves have size, the particles are sparse enough that the majority of the container's volume is essentially empty space compared to the size of the molecules.
However, it's important to note that this assumption is not universally true in all conditions. At very high pressures or low temperatures, the assumption of negligible molecular volume may not hold, and more sophisticated models may be required to accurately describe the behavior of the gas.