SCH4U Electron Configurations: Notes and Practice

Sodium explodes in water, neon does nothing at all, and chlorine grabs electrons from almost anything. The periodic table groups these behaviours into columns, but the reason behind the columns is how electrons are arranged in orbitals. An electron configuration is a shorthand that lists which orbitals an atom's electrons occupy. Once you can write one, you can explain an element's position in the table, the ions it forms, whether it is attracted to a magnet, and many of its trends.

Shells, subshells and orbitals. Electrons are organized in layers. A shell is all the orbitals with the same $n$ . Each shell is split into subshells labelled s, p, d and f, and each subshell contains a set number of orbitals. Since each orbital holds at most 2 electrons, the capacity of a subshell is twice its number of orbitals. Shell $n$ contains $n$ subshells, $n^{2}$ orbitals and at most $2 n^{2}$ electrons.

Subshell	Value of l	Number of orbitals	Maximum electrons	First appears in shell
s	0	1	2	n = 1
p	1	3	6	n = 2
d	2	5	10	n = 3
f	3	7	14	n = 4

Each subshell type has a fixed number of orbitals and a fixed electron capacity.

Three Rules for Filling Orbitals

Aufbau principle: electrons fill the lowest-energy orbitals available first, building up the atom one electron at a time.
Pauli exclusion principle: an orbital holds at most two electrons, and they must have opposite spins (drawn as ↑↓). No two electrons in an atom have the same four quantum numbers.
Hund's rule: when filling orbitals of equal energy (such as the three 2p orbitals), put one electron in each orbital with parallel spins before pairing any. Electrons repel, so they spread out first.

Fig 1Follow the diagonal arrows to get the filling order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p.

Why does 4s fill before 3d? In a multi-electron atom, subshells within a shell no longer have equal energy. An s electron spends part of its time very close to the nucleus, so it feels a stronger attraction and is lowered in energy. For potassium and calcium, the 4s orbital ends up lower than 3d, so it fills first.

Element	Full configuration	Orbital diagram of the last subshell	Unpaired electrons
C (Z = 6)	$1 s^{2} 2 s^{2} 2 p^{2}$	2p: ↑ \| ↑ \| _	2
N (Z = 7)	$1 s^{2} 2 s^{2} 2 p^{3}$	2p: ↑ \| ↑ \| ↑	3
O (Z = 8)	$1 s^{2} 2 s^{2} 2 p^{4}$	2p: ↑↓ \| ↑ \| ↑	2
Ne (Z = 10)	$1 s^{2} 2 s^{2} 2 p^{6}$	2p: ↑↓ \| ↑↓ \| ↑↓	0

Hund's rule in action: the 2p electrons spread out before they pair.

Noble-gas shorthand. Full configurations get long. Because a noble gas has a completely filled set of subshells, you can replace the inner electrons with the symbol of the previous noble gas in square brackets and write only what comes after. Iron (Z = 26) is $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{6} 4 s^{2} 3 d^{6},$ or simply $[Ar] 4 s^{2} 3 d^{6}$ . Some tables list subshells by shell number instead ( $[Ar] 3 d^{6} 4 s^{2}$ ); both describe the same atom. The electrons written after the bracket include the valence electrons, which control chemical behaviour.

The chromium and copper exceptions

The aufbau order predicts

[Ar] 4 s^{2} 3 d^{4}

for chromium and

[Ar] 4 s^{2} 3 d^{9}

for copper. The actual ground states are

[Cr] = [Ar] 4 s^{1} 3 d^{5}

and

[Cu] = [Ar] 4 s^{1} 3 d^{10}

. The 4s and 3d energies are very close, and a half-filled (

d^{5}

) or completely filled (

d^{10}

) d subshell is especially stable, so moving one electron from 4s to 3d lowers the atom's total energy. Molybdenum and silver, directly below them, show the same pattern.

Configurations of Ions

Anions: add electrons into the next available orbitals. $Cl$ is $[Ne] 3 s^{2} 3 p^{5}$ , so $Cl^{-}$ is $[Ne] 3 s^{2} 3 p^{6}$ , the same as argon.
Main-group cations: remove electrons from the outermost shell. $Na$ is $[Ne] 3 s^{1}$ , so $Na^{+}$ is $[Ne]$ .
Transition-metal cations: remove electrons from the highest shell first, so the 4s electrons leave before any 3d electrons, even though 4s filled first. $Fe^{2 +}$ is $[Ar] 3 d^{6}$ and $Fe^{3 +}$ is $[Ar] 3 d^{5}$ .
Atoms and ions with the same configuration are isoelectronic: $O^{2 -}$ , $F^{-}$ , $Ne$ , $Na^{+}$ and $Mg^{2 +}$ all have $1 s^{2} 2 s^{2} 2 p^{6}$ .

Blocks of the Periodic Table and Trends

Block	Groups	Subshell being filled	Typical properties
s	1 and 2 (plus He)	ns	Reactive metals with low ionization energies that form 1+ or 2+ ions
p	13 to 18	np	Nonmetals, metalloids and some metals; group 18 has filled s and p subshells and is very unreactive
d	3 to 12	(n − 1)d	Transition metals: often several possible ion charges, many coloured compounds, many used as catalysts
f	Lanthanides and actinides	(n − 2)f	Inner transition metals; the actinides are all radioactive

The shape of the periodic table comes directly from the capacities of s, p, d and f subshells: 2, 6, 10 and 14 columns wide.

The period number tells you the highest shell being filled, and the group tells you how many valence electrons there are. That explains the trends. Across a period, electrons are added to the same shell while the nuclear charge increases, so the outer electrons are pulled in more tightly: atomic radius decreases and ionization energy and electronegativity increase. Down a group, each new shell puts the valence electrons farther from the nucleus and behind more inner electrons that shield them, so radius increases and ionization energy decreases. Configurations even explain the small dips: boron's ionization energy is lower than beryllium's because its outer electron is in a higher-energy 2p orbital, and oxygen's is lower than nitrogen's because oxygen's fourth 2p electron must pair up and is repelled by its partner.

Unpaired electrons and magnets

Atoms or ions with unpaired electrons are paramagnetic: they are weakly attracted into a magnetic field. If every electron is paired, the substance is diamagnetic and is very weakly repelled. Counting unpaired electrons from an orbital diagram lets you predict which is which.

Electron Configurations

Key things to remember

Three Rules for Filling Orbitals

Configurations of Ions

Blocks of the Periodic Table and Trends

Worked examples

Where marks get lost

Flashcards 9

Practice