Electron configuration (intro)
The idea
Electrons around a nucleus are not scattered at random — they fill energy levels in a strict order, and writing that order down is the electron configuration. It matters because the outermost electrons decide everything about an element's bonding behavior, and because the periodic table's entire shape is a map of how those levels fill. You already know each element's atom has Z electrons when neutral; the configuration says where they sit.
Picture seats filling from lowest energy upward (the aufbau principle). The levels split into subshells with fixed capacities: any s subshell holds 2 electrons, p holds 6, d holds 10. The filling order starts 1s, 2s, 2p, 3s, 3p, 4s, 3d, and the superscript on each subshell counts the electrons placed there, so all superscripts must add up to the total electron count.
A common stumble is assuming the third shell must be completely full before the fourth starts. It is not: 4s sits lower in energy than 3d, so potassium and calcium fill 4s while 3d is still empty. The table itself encodes this — the row (period) an element sits in tells you its outermost occupied shell, and the column tells you how the outer subshells end.
Worked example
Write the full electron configuration of a neutral sulfur atom (Z = 16), state how many valence electrons it has, and predict the charge of the ion sulfur is most likely to form.
- Sulfur has 16 protons, so a neutral atom has 16 electrons to place in order of increasing energy.
- Fill the subshells in sequence: 1s² uses 2 (14 left), 2s² uses 2 (12 left), 2p⁶ uses 6 (6 left), 3s² uses 2 (4 left), and the final 4 go into 3p, giving 3p⁴.
- Write it out and audit the count: 1s² 2s² 2p⁶ 3s² 3p⁴, and 2 + 2 + 6 + 2 + 4 = 16, matching Z. The configuration is complete.
- Valence electrons are those in the outermost shell, n = 3 here: 3s² 3p⁴ gives 6 valence electrons. That is consistent with sulfur sitting in period 3 of the table, in the oxygen family.
- An atom with 6 valence electrons is 2 short of a stable octet, so sulfur tends to gain 2 electrons and form the S²⁻ ion rather than lose 6.
Answer. Sulfur is 1s² 2s² 2p⁶ 3s² 3p⁴, has 6 valence electrons, and typically forms S²⁻.
Check your understanding
- Why do elements in the same column of the periodic table show such similar chemistry, in terms of their configurations?
- How would the configuration of the S²⁻ ion differ from neutral sulfur, and which noble gas would it match?
- Why do electrons fill the lowest-energy subshells first, and what would be unstable about doing otherwise?
- How does an element's period number connect to its electron configuration, and what does its group tell you?
Build the foundations first
Electron configuration (intro) builds on these concepts. If any feel shaky, start there.