Ne 1s² 2s² 2p⁶

Be 1s² 2s²

Ga 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p¹

Sr²⁺ 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶

The best way to memorize electronic configurations, which are crucial to understanding atomic structure and reactivity, is to see the periodic table as being divided into regions corresponding to the s, p, d, and f orbitals. The row and region for each element tells you the quantum number for the last electron in the electronic configuration. Since all electronic configurations have the same sequence of orbitals, all you need to know is where to stop (at the last electron).

The four regions of the periodic table are:

s-orbitals: Columns 1A (alkali metals) and 2A (alkaline earth metals). The s-orbitals go from hydrogen, H (row 1), to francium, Fr (row 7), corresponding to principal quantum levels 1 through 7.

p-orbitals: Columns IIIA through VIIA (halogens). The p-orbitals (e.g., in Column IIIA) go from boron, B (row 2), to thallium, Tl (row 6), corresponding to principal quantum levels 2 through 6. There is no 1p orbital.

d-orbitals: Columns IIIB through IIB. These are the transition metals. The d-orbitals (e.g., in Column IIIB) go from scandium, Sc (row 4), to actinium, Ac (row 7), corresponding to principal quantum levels 3 through 6. There are no 1d or 2d orbitals. Each element in the first row of transition metals has a full 4s orbital and a partially full 3d orbital (zinc's 3d orbital is full, with 10 electrons).

f-orbitals: The lanthanides and actinides are the two rows at the bottom of the periodic table. They really are inserted into rows 6 and 7, after La and Ac, respectively. The last electron in Ce through Lu is a 4f electron. The last electron in Th through Lw is a 5f electron. There are no 1f, 2f, or 3f orbitals.

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