Why do 1s and 2s orbitals intersect? So I was going through the Schrödinger wave equation to determine the shapes of orbitals. We know that an orbital is nothing more than a region where probability of finding an electron is maximum.
Further, we know that the probability of finding a 2s electron in 1s orbital is just too low! Why do 1s orbital and 2s orbital intersect? I mean both the orbitals have the same shape and the radial function of 2s orbital is more than that of 1s which clearly means that 1s orbital lies inside 2s orbital!
But as it's impossible to find a 2s orbital electron in 1s orbital, shouldn't the shape of 2s orbital be something like a hollow sphere with thick walls instead of a complete sphere?
I read your question a few times, and I think you're confusing occupancy with spatial probability distribution.
Orbitals have a lot of spatial overlap. But if the orbital has two electrons “in it,” then it cannot describe the probability behavior of any other electrons in that atom, as other electrons must be in another orbital.
A 2s electron absolutely can go into a 1s orbital provided there is space. This is exactly what happens with atomic emission spectra when an electron in energy level $n=2$ transitions to $n=1$ and releases energy.