Please help, about real gas deviation from ideal behavior.

when we plot graph against P and PV, at 17¡ãC,:
for H2, there's straight ascending line.
for N2, line a little curved at low P then become ascending straight like H2.
for CO2, line curved, very descending at low P, then ascend but don't reach the ideal line.
please can you explain this in terms of P, V, T and inter molecular forces?

To understand the behavior of real gases, we need to consider their intermolecular forces. Intermolecular forces are the attractive forces between molecules that influence their behavior. The three types of intermolecular forces are London dispersion forces (present in all molecules), dipole-dipole forces (present in polar molecules), and hydrogen bonding (a special type of dipole-dipole force).

The ideal gas law, PV = nRT, assumes that the gas molecules have no volume and do not interact with each other. However, real gases deviate from this behavior due to intermolecular forces and the finite size of gas molecules.

Now, let's consider the behavior of each gas you mentioned:

1. Hydrogen (H2):
Hydrogen gas consists of diatomic molecules (H2). The intermolecular forces in hydrogen are predominantly London dispersion forces. These forces are weak and depend on the temporary fluctuations in electron distribution. Since hydrogen molecules are small and have low boiling points, the intermolecular forces are relatively weak. As a result, hydrogen gas behaves nearly ideal, and its graph against P and PV is a straight ascending line.

2. Nitrogen (N2):
Nitrogen gas also consists of diatomic molecules (N2). Similar to hydrogen, nitrogen predominantly exhibits London dispersion forces, but these forces are slightly stronger due to the larger size and greater electron density of nitrogen molecules. As a result, at low pressures, the curve in the graph indicates slight deviations from ideal behavior due to the increasing influence of intermolecular forces. However, at higher pressures, the intermolecular forces become less significant, and the graph becomes a straight line resembling ideal behavior.

3. Carbon dioxide (CO2):
Carbon dioxide is a linear triatomic molecule with two oxygen atoms bonded to a central carbon atom. CO2 molecules exhibit dipole-dipole forces due to the polarity of the molecule. The oxygen atoms are more electronegative than the carbon atom, creating a partial negative charge on the oxygen atoms and a partial positive charge on the carbon atom. These dipole-dipole forces are stronger than London dispersion forces. At low pressures, the graph shows a curved shape due to the stronger intermolecular forces, causing a decrease in volume compared to ideal gas behavior. As the pressure increases, the intermolecular forces become less significant, resulting in an ascending curve that does not reach the ideal line.

In summary, the behavior of real gases can be explained by considering the intermolecular forces present in the gas molecules. The strength of these forces and the molecular size play a significant role in determining the deviations of real gases from ideal gas behavior.