Here's why:

* Intermolecular forces: N2 is a nonpolar molecule, held together by weak London dispersion forces. HBr is a polar molecule, with a dipole moment, leading to stronger dipole-dipole interactions.

* Stronger intermolecular forces lead to deviations from ideal gas behavior. Ideal gases are assumed to have no intermolecular forces.

* Size and Polarizability: N2 is smaller and less polarizable than HBr.

* Larger size and higher polarizability contribute to stronger London dispersion forces. This again leads to deviations from ideal gas behavior.

Ideal Gas Behavior:

Ideal gases are theoretical entities that follow these assumptions:

* No intermolecular forces: The gas molecules are assumed to not attract or repel each other.

* Point masses: Molecules are assumed to have no volume.

* Elastic collisions: Collisions between molecules are perfectly elastic, conserving kinetic energy.

In summary: N2 is more ideal than HBr because it has weaker intermolecular forces and smaller size, making it closer to the assumptions of an ideal gas.