Is there any reason as to why we assume the volume of ideal gas to be negligible in comparison to the container? What is the underlying reason to assume that the ideal gas molecules are point molecules? Why can we not assume them to be non-point particles?
Because it gives simpler-to-derive laws which are often very good approximations
Clearly real gases do not always follow the ideal gas laws. They mostly liquefy under some conditions, for example, and, under those conditions they are clearly not ideal.
But in practice gas laws are used for things far away from those non-ideal regions. When we are applying the laws we are usually applying them to gases where the deviations from ideality are small. At normal lab conditions (1 atm pressure, room temperature) most of the gases you will ever deal with will be as close to ideal as matters within measurement error. Given that you want the simplest law possible and that law is easily derived by making the assumptions that particles have no size and don't stick together.
More complicated gas laws can and have been developed (van der Waals equation was a 19th century development of a simple extended law that captures some features about liquefaction and non-ideality under cold or high pressure conditions). But the deviations of the results from the simplest ideal gas equation are small under normal conditions (van der Waals calculations for carbon dioxide differ from ideal results by 0.5% under normal lab conditions). This is such a small amount that more complex equations are not worth bothering with unless you know you are using very low temperatures or high pressures.
And, when you are in one of those regions, you will find that the gas equations get very complicated very quickly. It takes a lot of extra empirical data and mathematical complexity to handle those cases and under normal conditions this just isn't worth the effort. Why waste effort when you don't gain any practical benefit?