TL;DR:
An electron configuration lists how many electrons sit in each subshell, such as oxygen's
1s² 2s² 2p⁴.Subshells hold s: 2, p: 6, d: 10, f: 14 electrons, and fill roughly in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d.
The table's blocks and rows mirror the filling order, so you can read a configuration off the table.
The periodic table is not arranged by accident. Its rows and columns come from how electrons fill the space around a nucleus. Once you see that, the table stops being something to memorise and becomes something you can read.
Shells, subshells and orbitals
Electrons live in shells numbered 1, 2, 3 and so on, with higher numbers farther from the nucleus. Each shell contains subshells named s, p, d and f. A subshell is made of orbitals, and each orbital holds at most 2 electrons.
Subshell | Orbitals | Maximum electrons |
|---|---|---|
s | 1 | 2 |
p | 3 | 6 |
d | 5 | 10 |
f | 7 | 14 |
A configuration is written as shell number, subshell letter and electron count as a superscript, so 2p⁴ means four electrons in the 2p subshell.
The filling order
Electrons fill the lowest-energy subshell that still has room. This is the Aufbau principle. The order for most atoms is:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6pNotice that 4s fills before 3d. A subshell's energy depends on more than its shell number. Two further rules apply: Pauli exclusion says an orbital holds at most two electrons with opposite spins, and Hund's rule says electrons in a subshell spread into separate orbitals before pairing up.
Worked example: oxygen
Oxygen has atomic number 8, so a neutral atom has 8 electrons.
Fill 1s: 2 electrons. Remaining: 6.
Fill 2s: 2 electrons. Remaining: 4.
Put the last 4 in 2p, which can hold 6. Remaining: 0.
Write it:
1s² 2s² 2p⁴.
Check the total: 2 + 2 + 4 = 8, matching the atomic number. The outer shell is n = 2, with 2 + 4 = 6 valence electrons, which is why oxygen forms two bonds, needing two more to reach eight.
Hund's rule decides how the four 2p electrons sit: one pair in one orbital and one electron in each of the other two, so there are two unpaired electrons. That fits oxygen's tendency to form two bonds.
More examples
Chlorine (17):
1s² 2s² 2p⁶ 3s² 3p⁵. Check:2 + 2 + 6 + 2 + 5 = 17. Seven valence electrons.Sodium (11):
1s² 2s² 2p⁶ 3s¹, or[Ne] 3s¹in noble-gas shorthand. One valence electron, which it loses easily to become Na⁺.Iron (26):
[Ar] 4s² 3d⁶. Check: argon has 18, and18 + 2 + 6 = 26.
Reading the periodic table as a map
The table splits into blocks named after the subshell being filled.
s-block: the first two columns, filling s subshells (2 columns for 2 electrons).
p-block: the right-hand six columns, filling p subshells (6 columns for 6 electrons).
d-block: the transition metals in the middle, ten columns for 10 electrons.
f-block: the two rows set apart at the bottom, fourteen columns for 14 electrons.
For the main-group elements, the column number tells you the valence electrons: sodium (column 1) has 1, chlorine (column 17) has 7, oxygen (column 16) has 6. The row number is the outer shell number. Oxygen is in row 2, so its outer shell is n = 2. Elements in the same column share an outer configuration, which is why they behave alike.
A few exceptions: Chromium and copper do not follow the simple order. Chromium is
[Ar] 4s¹ 3d⁵and copper is[Ar] 4s¹ 3d¹⁰, because half-filled and full d subshells are especially stable. Your course will list which exceptions you need.
Ions follow the same bookkeeping. Sodium becomes Na⁺ by losing its 3s electron, leaving 1s² 2s² 2p⁶, the same as neon. Chlorine becomes Cl⁻ by gaining one electron into 3p, giving 1s² 2s² 2p⁶ 3s² 3p⁶, the same as argon. Atoms form the ions that give them a noble-gas configuration, which is why sodium chloride is Na⁺Cl⁻ and not something else.
Here is a quick table-reading exercise for sulfur, atomic number 16. It is in row 3, so its outer shell is n = 3. It is in column 16 of the p-block, so it has 6 valence electrons. Fill the order: 1s² 2s² 2p⁶ 3s² 3p⁴. The total is 2 + 2 + 6 + 2 + 4 = 16. Sulfur sits directly under oxygen and has the same outer pattern, s² p⁴, which is why the two behave alike.
How to check your answer
Sum. The superscripts must add up to the atomic number (minus the charge for an ion).
Capacity. No subshell exceeds its maximum: 2, 6, 10 or 14.
Valence count. The electrons in the highest shell should match the group.
Ions. Remove electrons from the highest shell first, not the last one added.
Mistakes that cost marks
Filling 3d before 4s.
Putting 8 electrons in a 2p subshell.
Counting the d electrons of a transition metal as valence electrons without checking the course convention.
Forgetting to subtract or add electrons for ions.
Mixing the shell number and the superscript, writing
2p⁴as 2 times 4.
Where you will use it next
Valence electrons feed directly into bonding and shapes, as in predicting molecule shapes with VSEPR. They also explain trends across the table in periodic trends: atomic radius and electronegativity.
Frequently asked questions
Why does 4s fill before 3d?
In a many-electron atom, a 4s electron has a lower energy than a 3d one at that point. Once the 3d subshell has electrons, the energies shift, which is why 4s electrons leave first when a transition metal forms an ion.
What does [Ne] mean in a configuration?
It stands for neon's configuration, 1s² 2s² 2p⁶, so you only write the electrons beyond it.
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