The periodic table: classification of elements

By **increasing atomic number** (number of protons), not by relative atomic mass.

A horizontal **row**; the period number equals the highest occupied **main energy level (n)**.

A vertical **column**; elements in a group have the **same number of outer (valence) electrons**.

The **sublevel** that the outermost electrons are filling (s, p, d or f).

Groups **1 and 2** (plus H and He) — outer electrons fill the **s** sublevel.

Groups **13–18** — outer electrons fill the **p** sublevel.

The **centre** of the table (groups 3–12) — the **transition metals**, filling the d sublevel.

The **two detached rows** at the bottom — the **lanthanides and actinides**, filling the f sublevel.

Name the **sublevel the outermost electron enters** (e.g. …3p⁵ → p-block; …3d⁶ → d-block).

**Period** number = n of the outer shell; **group** number = number of outer electrons (group 17 → 7).

The **s-block** — its next electron would enter the **8s** sublevel (group 1, period 8).

**Nuclear charge** (proton pull) and **shielding/distance** (inner shells + extra shells).

The energy needed to remove one mole of electrons from one mole of **gaseous** atoms: X(g) → X⁺(g) + e⁻.

**Half** the distance between the nuclei of two bonded atoms — a measure of atom size.

How strongly an atom attracts a **bonding pair** of electrons (Pauling scale).

The energy change when one mole of gaseous atoms **gains** an electron: X(g) + e⁻ → X⁻(g).

**Decreases** — greater nuclear charge with similar shielding pulls the outer shell in.

**Increases** — each element has an extra electron shell.

**Increases** across a period (stronger pull); **decreases** down a group (further out, more shielded).

**Increases** across a period, **decreases** down a group (fluorine is the most electronegative).

A cation is **smaller** than its atom (it often loses a whole shell).

An anion is **larger** than its atom (extra electron–electron repulsion spreads the shell out).

Compare **nuclear charge**, compare **shielding/distance**, then state the **net effect** (held more/less tightly).

The same number of **outer (valence) electrons**, so they react in similar ways.

It **increases** — the outer electron is further out and more shielded, so it is **lost more easily**.

It **decreases** — the atom is bigger, so an incoming electron is **harder to gain**.

K is lower in group 1: **bigger atom + more shielding** → outer electron lost more easily.

F is smaller with less shielding, so it **gains** an electron more easily.

Able to act as **both an acid and a base** — reacts with acids **and** alkalis (e.g. Al_{2}O_{3}).

It **decreases** — elements change from **metallic** (Na) to **non-metallic** (Cl, Ar).

**Basic → amphoteric → acidic** left to right (Na_{2}O/MgO basic, Al_{2}O_{3} amphoteric, SO_{3} acidic).

**Basic** (e.g. Na_{2}O, MgO). Non-metal oxides are **acidic** (e.g. SO_{3}, P_{4}O_{10}).

**Caesium + fluorine** — lowest (most reactive) metal + top (most reactive) halogen.

Li < Na < K < Rb < Cs (increases down).

F > Cl > Br > I (decreases down).

A **d-block metal** that forms **at least one stable ion with a partially filled d sub-shell**.

Their only ions are **Sc³⁺ ([Ar] 3d⁰)** and **Zn²⁺ ([Ar] 3d¹⁰)** — empty/full d, never **partially filled**.

**4s fills first** (slightly lower energy when empty), so atoms end in **3d^{x} 4s²**.

**Remove 4s electrons before 3d.** e.g. Fe²⁺ = [Ar] 3d⁶ (the two 4s electrons go first).

**[Ar] 3d⁵ 4s¹** — an anomaly; a **half-full 3d⁵** is extra stable.

**[Ar] 3d¹⁰ 4s¹** — an anomaly; a **full 3d¹⁰** is extra stable.

The **4s and 3d sub-shells are close in energy**, so electrons can be removed in steps for similar energies → several stable states.

**+2 and +3** (Fe²⁺ = [Ar] 3d⁶; Fe³⁺ = [Ar] 3d⁵, a stable half-full sub-shell).

**+1 and +2** (Cu⁺ in Cu_{2}O, Cu²⁺ in CuSO_{4}).

**+7** — four O at −2 (−8) with an overall −1 charge forces Mn to +7.

The **partially filled d sub-shell** splits in a ligand field and **absorbs visible light**.

They can **change oxidation state** and use empty/part-full **d orbitals** to bind reactants (e.g. Fe in the Haber process).

Having **unpaired d electrons** — these are drawn into a magnetic field.

The ligands **split** the five d orbitals into two energy levels separated by a gap **Δ** (the splitting energy).

A molecule or ion (e.g. H_{2}O, NH_{3}, CN⁻) that bonds to a central metal ion by donating a **lone pair** of electrons.

The **splitting energy** — the energy gap between the two split d-orbital levels in a complex.

A d electron **absorbing a photon** of energy equal to Δ and jumping from the lower d-orbital level to the upper one.

Δ matches the energy of **visible light**, so the complex absorbs part of the visible spectrum in a d-d transition.

The **complement** of the colour **absorbed** — the leftover light. Absorbs red → looks green; absorbs blue → looks orange.

The **metal + its oxidation state**, the **ligand** (spectrochemical series), and the **number/geometry** of ligands.

Ligands ranked by the size of Δ they cause: I⁻ < Cl⁻ < H_{2}O < NH_{3} < CN⁻ (weak-field → strong-field, small Δ → large Δ).

**Weak-field** (I⁻, Cl⁻) → **small** Δ; **strong-field** (CN⁻, CO) → **large** Δ.

Larger Δ needs a **higher-energy** photon → a **shorter** wavelength of light is absorbed.

A different oxidation state changes Δ, so a **different wavelength** is absorbed and the complementary colour seen changes.

A **tetrahedral** complex has a **smaller Δ** than the equivalent **octahedral** one, so it absorbs different light and shows a different colour.