Atomic Structure and Electron Configuration

Electrons are arranged in energy levels (shells), which are further divided into subshells, and each subshell contains orbitals.

The Four Types of Subshell

Each orbital holds a maximum of 2 electrons with opposite spins (Pauli exclusion principle).

Energy Level and Subshell Structure

Filling Order

Subshells fill in order of increasing energy (not always by shell number):

$1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d \rightarrow 4p \rightarrow 5s \rightarrow 4d \rightarrow 5p$

Key point: 4s fills before 3d because 4s is lower in energy.

Aufbau Principle

Electrons fill orbitals starting from the lowest energy level available.

Pauli Exclusion Principle

Each orbital holds a maximum of 2 electrons with opposite spins (↑↓).

Hund's Rule

Electrons occupy orbitals in a subshell singly first (with parallel spins) before pairing up.

For example, the 2p subshell with 3 electrons: ↑ ↑ ↑ (one in each orbital) With 4 electrons: ↑↓ ↑ ↑ (pair in first, single in remaining two)

Notation

Full notation shows each subshell and the number of electrons:

Exceptions

Chromium (Cr, Z=24): Expected $[\text{Ar}] \, 3d^4 \, 4s^2$, actual: $[\text{Ar}] \, 3d^5 \, 4s^1$

Copper (Cu, Z=29): Expected $[\text{Ar}] \, 3d^9 \, 4s^2$, actual: $[\text{Ar}] \, 3d^{10} \, 4s^1$

Reason: Half-filled and fully-filled d subshells are particularly stable.

Ion Configurations

When forming ions, 4s electrons are lost before 3d electrons.

• Fe: $[\text{Ar}] \, 3d^6 \, 4s^2$
• Fe²⁺: $[\text{Ar}] \, 3d^6$ (loses 2 × 4s electrons)
• Fe³⁺: $[\text{Ar}] \, 3d^5$ (loses 2 × 4s + 1 × 3d)

The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms:

$\text{X}(g) \rightarrow \text{X}^+(g) + e^-$

The second ionisation energy: $\text{X}^+(g) \rightarrow \text{X}^{2+}(g) + e^-$

Successive ionisation energies always increase because:

• Each electron is removed from an increasingly positive ion
• Greater nuclear attraction must be overcome

Factors Affecting Ionisation Energy

1. Nuclear charge: More protons → stronger attraction → higher IE
2. Atomic radius: Further from nucleus → weaker attraction → lower IE
3. Shielding: More inner electrons → more shielding → lower IE
4. Subshell: Paired electrons are slightly easier to remove (electron-electron repulsion)

Trends

Across a period: IE generally increases (more protons, same shielding)

Down a group: IE decreases (more shells, greater distance, more shielding)

Anomalies in Period 3

1. Be → B drop: B has $2s^2 \, 2p^1$; the 2p electron is easier to remove than a 2s electron (2p is higher energy, less penetrating)

2. N → O drop: O has $2s^2 \, 2p^4$; the extra electron in the 2p is paired, and electron-electron repulsion makes it easier to remove

A graph of successive IEs for an element shows jumps that indicate the electron shell structure.

For sodium ($1s^2 \, 2s^2 \, 2p^6 \, 3s^1$):

• 1st IE: relatively low (3s¹ electron)
• Big jump to 2nd IE (removing from 2p⁶ — inner shell, closer to nucleus)
• IEs 2–9 increase gradually (removing 2s²2p⁶)
• Another big jump at 10th IE (removing from 1s²)

The pattern of jumps tells you the number of electrons in each shell.

When atoms are excited (heated or given electrical energy), electrons jump to higher energy levels. When they fall back, they emit photons of specific energies.

$\Delta E = h\nu = \frac{hc}{\lambda}$

where $h = 6.63 \times 10^{-34}$ J·s, $c = 3.00 \times 10^8$ m/s

The line spectrum of an element (discrete lines at specific wavelengths) provides evidence for quantised energy levels.

• Always write configurations using the subshell notation ($1s^2 \, 2s^2$ etc.)
• For ions: remove 4s electrons before 3d electrons
• Know the Cr and Cu exceptions and be able to explain them
• In IE questions, identify where the big jumps are
• For anomalies across a period, mention subshell energy (s vs p) and electron pairing

• Electrons occupy subshells (s, p, d, f) within energy levels
• Filling order: $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p...$
• Aufbau, Pauli, and Hund's rule govern filling
• Ionisation energy increases across a period (with anomalies) and decreases down a group
• Successive IE jumps reveal electron shell structure
• Cr and Cu are exceptions ($3d^5 4s^1$ and $3d^{10} 4s^1$)