The number of electrons an atom has is equal to its atomic number. A hydrogen atom has 1 electron, a carbon atom has 6, sulfur has 16, iron 26, gold 79, and uranium 92. Those electrons can't just wander anywhere inside the atom — each one is assigned a region of space it can occupy, a shape called an orbital. This module looks at what orbitals exist and what they look like.

Electron orbitals are grouped into shells, numbered from 1 upward. This number, n, is the principal quantum number, and the larger it is, the larger the orbital.

Each shell has one or more subshells. Each subshell contains orbitals of similar shape. The subshell types used by real atoms are s, p, d, and f. As the shell number increases, the number of subshells in it also increases:

Shell	Subshells

Shells don't stop at 6 — they keep going up, gaining more subshell types each time. But real atoms only use shells through n = 7, and only the subshells shown below:

Shell	Subshells used by real atoms

The shapes of the orbitals within the first three shells are shown below.

• s — spherical, getting larger as n increases
• p — dumbbell-shaped, getting larger as n increases
• d — cloverleaf, except for the 3d_z² orbital, which looks like a p orbital with a ring around its waist
• f — three-dimensional shapes more complex than d orbitals

Subshell	# of orbitals	Orbitals
3s	1	3s
3p	3	3px  3py  3pz
3d	5	3dz2  3dxz  3dyz  3dxy  3dx2−y2

Shell	Subshells (# of orbitals)

Now it's time to practice. There are three levels: which subshells exist, which orbital names are valid, and what the orbital shapes look like. Each level has a reference sheet you can pop open at any time.

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