Order of the subshells in an electron configuration According to what my chemistry teacher said, the energy of a subshell is given by the sum of its principal quantum number (n) and its second quantum number (l).

e.g.:


1s = n+l = 1+0 = 1
2s = n+l = 2+0 = 2
2p = n+l = 2+1 = 3
3s = n+l = 3+0 = 3


However, when it comes to make a graphical representation of an electron configuration (which follows an order of increasing energy), she doesn't put the subshells having the same energy in the same level, like in the following picture:



Shouldn't the subshells 2p and 3s (which both have an energy equals to 3), as well as 4p and 3d, be put in the same level, since they have the same energy?
Whilst the said concept is correct, as orbitals increase in number the less they are attracted to the centre of nuclei, consequently the further they move away the higher the energy they possess.

The 2s electrons spent more time closer to the nucleus (when compared to 2p electrons) and this reduces the energy of electrons in s orbitals, the nearer they get, the lower the energy. This is why the s orbital always has a slightly lower energy than p orbital at the same energy level.

The 2s, 3s .... orbitals get progressively away from the nucleus.