Teaching transition metals and complex ions

Why is blood red? Why, for some sea creatures, is blood blue? The answer lies in the fascinating and colourful topic of transition metal complexes. While the coordination of iron(II) is responsible for the red colour in haemoglobin, it is copper(II) that gives haemocyanin its blue colour. When deoxygenated, it is colourless; when oxygenated, it’s blue. Transition metal ions, surrounded by ligands coordinating to a central metal ion, have been used in medicine for decades – in applications such as anti-cancer treatments or as tracers providing diagnostic information.

D-orbitals, electrons and energy gaps can seem hugely abstract to students, so real-world contexts, and the wonderfully colourful chemistry of transition metal complexes – that they can actually see – offer a tangible route to learning the underlying chemistry.

What students need to know

To understand transition metal ions and their complexes, students need to be secure in their grasp of detailed electronic arrangements and d-subshells. A transition metal complex consists of a central metal ion surrounded by ligands. A ligand is an ion or molecule with a lone pair of electrons to donate to form a coordinate (dative covalent) bond with a transition metal ion. The coordination number of the complex is the total number of bonds from the ligands to the central transition metal.

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Familiarise learners with example ligands to secure understanding of the complexes they form

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Link colour to d-orbital splitting with a clear example of how orbitals split in a complex

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The d-orbitals are no longer degenerate in a complex of a transition metal due to the repulsions of orbitals on particular axes with approaching ligands and a change in energy.The colours of most transition metal complexes are due to d−d electron transitions absorbing the complementary colour of light to what’s observed. The colour depends on the transition metal, oxidation state, ligands and coordination number. A transition metal can exist in a range of oxidation states in its compounds, and can be readily inter-converted.

Common misconceptions

• The word complex can be misleading. Explain to your students that it doesn’t mean complicated in this context, and refers to the compound consisting of multiple parts.
• Electronic configurations with filling of d-orbitals, sub-shells and d−d transition of electrons can seem abstract to students.
• Learners can confuse absorption and emission. The absorption of light will show complementary colours, eg a solution appears violet in colour because the solution absorbs yellow light, which is the complementary colour to violet.
• Students can find it hard to fully explain why zinc compounds are colourless.
• They need to understand that both electrons in the coordinate (dative covalent) bond come from the same atom.
• Learners can also struggle to recognise a change in oxidation state.
• Many of the terms and vocabulary used to teach this topic will be new to students. They can find it hard to name complexes due to their length and pupils’ lack of familiarity with some of the parts.

Ideas for your classroom

When introducing transition metal ion complexes, start with a simple, familiar example. A test tube of copper(II) sulfate in aqueous solution shows the well-known blue colour due to the Cu²⁺ ions complexing with water. Get students to draw this hexaaqua octahedral structure and write out the d-electrons and splitting. Add ammonia solution until the sample turns dark blue, as water ligands are replaced by ammonia, to show how a stronger field ligand affects the colour. Further outline the sequential replacement of ligands by looking at copper complexes. This can be used at a later stage to explore ligands, chelation and linking equilibria and stability constants.

Use microscale to give your students the opportunity to closely observe a selection of transition metal complexes, their colours, and the changes they undergo when adding other ligands and reagents. It also offers opportunities for discussion about the observed changes. Can your students identify evidence of complex formation? Is there a change in oxidation state? And why is there no colour with zinc?

Looking at vanadium complexes in more depth will prompt your students to consider variable oxidation states. Start from the yellow V(V) ions (VO₂⁺) that are reduced by zinc in an acidic solution first to the blue V(IV), and so on, all the way to the purple V(II) ions. Then use potassium permanganate (another transition metal with a high oxidation state) to oxidise vanadium back through all these changes to +5. Get your students to think about electrode potentials and making predictions.

Assessing student understanding

Throughout this topic, give students plenty of opportunity to practise the various key aspects, such as drawing d-electrons, showing what happens to electrons when energy is absorbed, and drawing complex ions showing the coordination of ligands with 3D geometry. If they answer on mini whiteboards, you can quickly view answers to assess who has understood or needs more help. Follow up with microscale practical activities so students can make links to concrete examples of selected metal ions, ligands and colours.