Coloured complexes and d-orbital splitting (HL)

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

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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.

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

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

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

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

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

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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).

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

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

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

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