It
doesn't include anything beyond n=4, and the energy differences are not to
scale; only the ordering is correct.
Due to the Pauli Exclusion Principle, each orbital, sublevel and primary level can only hold a certain number of electrons. Thus, if we thnk about adding electrons to a bare nucleus until we have the correct number for that atom, the electrons must go into a distinct set of orbitals. In the ground state of the atom, electrons will occupy the lowest energy orbitals first, and only fill the higher energy orbitals when no lower energy orbitals are left. The arrangement of electrons in an atom is known as its electronic configuration
A schematic diagram of the electron energy levels is at the right. It
doesn't include anything beyond n=4, and the energy differences are not to
scale; only the ordering is correct.
An an example, if we want to fill a magnesium atom with electron, we must place 12 electrons into various orbitals. We start by placing 2 electrons into the 1s orbital, which is its maximum capacity. The next lowest orbital is the 2s, which can also hold two electrons. We've now placed four electrons, eight to go.
The next sublevel up is the 2p. Since it has l=1 and thus three orbitals, it can hold six electrons. Filling it leaves us with two electrons. The next level is the 3s, which can hold two electrons. All electrons have now been placed, and we can write the electron configuration of the Mg atom as
Writing all this out can be tedious for an atom like uranium with 92 electrons. Thus, it is common to abbreviate the electron configuration by that of the preceding noble gas. For Mg, this is neon, which has an electron configuration of 1s22s22p6, so we can write the the electron configuration of the Mg atom as
Deciding how electrons distribute themselves in a given set of orbitals is the function of Hund's rule
Example: What is the ground state electron configuration of the iron atom?
Solution: The iron atom has 26 electrons.
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