Final answer:
Gases generally follow the ideal gas law because they have negligible volume compared to the container, which is most consistent with the properties of an ideal gas under low pressure and high temperature where intermolecular attractions are also negligible.
Step-by-step explanation:
The best explanation for why gases generally follow the ideal gas law is (c) Gases have negligible volume compared to the container. The ideal gas law assumes that the particles of a gas have no significant volume and do not attract or repel each other, which allows the particles to move about randomly and occupy the whole volume of their container. At low pressures and high temperatures, real gases behave closely to this model because the volume of the gas molecules is small compared to the container, and intermolecular attractions are negligible. Therefore, under such conditions, real gases approximate the behavior of an ideal gas, making the ideal gas law a good approximation for describing their properties.
In contrast, options a) and d) are incorrect because gases do have intermolecular forces, though they are weak, and gases are definitely composed of individual molecules. Option b) is also incorrect because gases do not have definite shape and volume; they take the shape and volume of their container.