Electron configurations

A tiny **packet of light energy**; its energy is given by E = hf (higher frequency → more energy).

A **fixed, allowed energy** an electron can have in an atom; energy levels are **discrete (quantised)**.

Continuous = an **unbroken rainbow** (all wavelengths). Line = a few **discrete bright lines** on black, from an excited element.

An excited electron **falls** from a higher to a lower energy level, emitting a photon of fixed energy (one line per allowed jump).

That the electron's energy levels are **discrete (quantised)** — fixed lines mean only fixed energy gaps are allowed.

The spectral lines get **closer together** toward **high frequency/energy**, because the energy levels bunch up at higher n.

The **biggest energy gap** — an electron falling **to n = 1** (the ground state).

$c = \lambda f$ — speed of light = wavelength × frequency, so **high f means short λ**.

$E = hf$ — photon energy = Planck's constant × frequency (higher f → higher E).

**Radio < red < violet** in frequency, so radio is lowest energy and violet is highest.

The lines merge; the electron gains just enough energy to **leave the atom** — this gives the **ionisation energy**.

The major 'shell' of an atom (n = 1, 2, 3, …); higher n means **higher energy** and **further** from the nucleus.

A subdivision of a main level, labelled **s, p, d, f**, differing slightly in energy (s < p < d < f).

A region around the nucleus that can hold up to **2 electrons**.

A **sphere** centred on the nucleus.

A **dumbbell** — two lobes pointing in opposite directions through the nucleus.

s = **1**, p = **3**, d = **5**, f = **7** orbitals.

s = **2**, p = **6**, d = **10**, f = **14** (2 electrons per orbital).

**2n²** — so n = 1 → 2, n = 2 → 8, n = 3 → 18, n = 4 → 32.

**4s** fills first — it is slightly lower in energy than 3d.

Electrons occupy orbitals of a sublevel **singly** (parallel spins) before any pairing up.

**s < p < d < f** (s is lowest, f is highest).

**3s, 3p and 3d** (max 18 electrons).

Electrons fill the **lowest-energy** sub-shell available first (build up: 1s, 2s, 2p, 3s, …).

Each orbital holds **at most 2 electrons**, and they must have **opposite spins**.

Within a sub-shell, electrons occupy orbitals **singly with parallel spins** before any pairing occurs.

1s, 2s, 2p, 3s, 3p, **4s, 3d**, 4p — note **4s fills before 3d**.

**s = 2**, **p = 6**, **d = 10** (each orbital holds 2).

1s² 2s² 2p⁶ 3s² 3p⁴.

Replace the inner electrons with the **previous noble gas** in [ ], then list the outer electrons — e.g. Ca = [Ar] 4s².

Start from the atom and **remove electrons from the highest main shell (largest n) first** — for transition metals, **4s before 3d**.

**[Ar] 3d⁶** — the two **4s** electrons are removed first, not the 3d.

**Add** the gained electrons to the next available sub-shell — e.g. O²⁻ = 1s² 2s² 2p⁶.

A **half-full** 3d⁵ sub-shell is extra stable, so one 4s electron promotes to 3d.

A **full** 3d¹⁰ sub-shell is extra stable, so one 4s electron promotes to 3d.