ENERGY LEVELS
Quantum theory has taught us that electrons in atoms are in discrete energy levels around the
nucleus. Immediate occupancy is in electron orbitals which are defined as regions of highest
probability of finding an electron. Originally, orbitals were viewed as non-discrete negatively
charged clouds. Quantum theory, however, has redefined atomic orbitals in terms of their size,
shape, and orientation from the nucleus. Each of these states is given by a separate quantum
number. To be complete there must be two electrons of opposite spin filling each orbital. The chart
below illustrates the energy levels and the type of orbitals that are allowed at each level. Electrons
fill vacant orbitals sequentially beginning with the lowest and going to the highest until the quota
represented by Z, the atomic number is met.
Principal
Quantum No. Energy
(n) Level Subshell(s) Filled configuration No. Electrons
1
First (K shell)
1s
1s2 2
2
Second (L shell)
2s 2p
2s22p6 8
3
Third (M shell)
3s 3p 3d
3s23p63d10 18
4
Fourth (N shell)
4s 4p 4d 4f
4s24p64d104f14 32
Rules Governing the Assignment of Electrons to Orbitals
1. Pauli exclusion principle. No two electrons in the same atom share the same set of quantum
numbers. Because of this rule, the two electrons that occupy orbitals of the same energy and
orientation must be of opposite spin, designated spin up and spin down. The spin state, therefore,
becomes a quantum number.
2. Hund’s rule of ground state electron configuration. For any set of orbitals of the same energy (p,
d, f) the lowest energy configuration is achieved by placing the electrons is different orbitals of the
set and with parallel spins. This rule recognizes that segregating electrons is a desired feature to
achieve a low energy state and that pairing electron is a higher energy state. We will see the
importance of this rule when we discuss 3d orbital splitting. You may also recall from organic
chemistry that one of the ground state configurations of carbon is designated sp3, meaning one
electron each in the 2s, 2px, 2py, and 2pz orbitals. This configuration gives rise to the familiar
tetrahedral structure of carbon compounds.
3. In the table above you will note that each spherical s orbital holds a maximum of two electrons,
each p orbital (better spoken of as a p set because the p set consists of three sausage shaped orbitals
of equal energy oriented along the x, y, and z axis) holds 6 electrons when filled, and each d set of
orbitals fills out at 10.
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