Electron configuration and the 2p orbital I looked through the questions already asked about this topic and they were very specific and over my head. I'm having trouble with the basics.

When drawing orbital diagrams I know electrons are placed in the 2p boxes singly first then paired because of they have parallel rotations and the Pauli exclusion principle states that an orbital can only house 2 electrons and those electrons must be of opposite spin. My question is why do electrons behave this way in the 2p orbital but not the 1s or 2s? And why does the p subshell contain three 2p orbitals?

As far as electron configurations and orbital diagrams are concerned I've gotten to Ne and the 2p configuration confused me so I though I'd ask before moving on.
first of all electrons do not rotate but they have spins
next magnetic quantum number decides how many electrons an orbital can have
s can have 2(m=1){consider 1 box}
p can have 6(m=3){consider 3 boxes}
d can have 10(m=5){consider 5 boxes}
f can have 14(m=7){consider 7 boxes}
where each box can have at max 2 electrons and each box represents different orientation of the orbital
and all the orbitals(except s) have degenerate orbitals, i.e they have the same energy so all have to filled with atleast one electron before pairing up. this rule is known as Hund's Rule of Maximum Multiplicity
there are are some more reasons fro it. they are
to maximize exchange energy 
to attain symmetry