Transition metals are the d-block elements in the central section of the periodic table, from titanium to copper and beyond. They are harder, denser, and higher-melting than the Group 1 alkali metals; they form coloured compounds; they can act as catalysts; and they commonly exist in more than one oxidation state.
Where are transition metals in the periodic table?
The transition metals occupy the d-block, the ten-column central section between Group 2 (alkaline earth metals) and Group 3 (aluminium group). In Period 4, these are the elements from titanium (Ti, atomic number 22) to copper (Cu, 29). Zinc (Zn, 30) is often included in discussions but is technically not a typical transition metal because its d-orbitals are completely filled in all its compounds.
Familiar transition metals include iron (Fe), copper (Cu), nickel (Ni), chromium (Cr), manganese (Mn), cobalt (Co), titanium (Ti), and tungsten (W). Together they account for many of the most industrially important metallic elements on Earth.
What physical properties do transition metals share?
Compared with the Group 1 alkali metals (lithium, sodium, potassium), transition metals have dramatically different physical properties:
| Property | Group 1 alkali metals | Transition metals |
|---|---|---|
| Density | Low (Li floats on water) | High (iron 7,900 kg/m³; copper 8,900 kg/m³) |
| Melting point | Low (Na melts at 98 °C) | High (iron 1,538 °C; tungsten 3,422 °C) |
| Hardness | Soft (cut with a knife) | Hard (structural materials) |
| Reactivity with water | Violent (Na explodes) | Slow or negligible (iron rusts; copper does not react) |
| Electrical conductivity | Good | Good to excellent (copper is the standard) |
These properties arise because transition metal atoms have more electrons available for metallic bonding — the delocalised electrons hold the lattice together with greater force — giving stronger, denser structures.
Why do transition metals form coloured compounds?
Transition metal compounds are characteristically coloured because their partially filled d-orbitals allow electrons to absorb specific wavelengths of visible light. When a photon of the right energy is absorbed, an electron jumps from a lower d-orbital to a higher one. The colour we see is the complementary colour of what was absorbed.
Common transition metal ion colours in aqueous solution:
| Ion | Colour | Example compound |
|---|---|---|
| Cu²⁺ | Blue | Copper(II) sulfate solution |
| Fe²⁺ | Pale green | Iron(II) chloride solution |
| Fe³⁺ | Orange-brown | Iron(III) chloride solution |
| Mn²⁺ | Pale pink | Manganese(II) sulfate solution |
| Cr³⁺ | Dark green | Chromium(III) chloride solution |
| Ni²⁺ | Green | Nickel(II) sulfate solution |
Group 1 and Group 2 metals form white or colourless compounds because their ions have no partially filled d-orbitals — there are no d-electron transitions available to absorb visible light.
Why can transition metals have variable oxidation states?
An oxidation state (or oxidation number) describes how many electrons an element has lost or gained in a compound. Most Group 1 metals have only one oxidation state (+1); Group 2 only +2. Transition metals commonly show two or more oxidation states because their d-electrons and their outer (4s) electrons are close in energy — different numbers can be removed depending on the conditions.
Iron, for example, forms both Fe²⁺ (iron(II), losing 2 electrons) and Fe³⁺ (iron(III), losing 3 electrons). The oxidation state affects the compound's colour, magnetic properties, and reactivity. In naming conventions, the Roman numeral in brackets tells you the oxidation state: iron(II) chloride = FeCl₂; iron(III) chloride = FeCl₃.
How do transition metals act as catalysts?
Transition metals and their compounds are excellent catalysts because their variable oxidation states allow them to accept and then donate electrons in a reaction cycle — effectively providing a lower-energy reaction pathway:
- A reactant molecule bonds to the transition metal surface or ion, temporarily changing the metal's oxidation state.
- The bonding weakens the bonds within the reactant, making it more reactive.
- The product forms and leaves the metal, which returns to its original oxidation state — ready to catalyse the next cycle.
Key industrial examples:
- Iron (Fe): catalyst in the Haber process for ammonia synthesis (N₂ + 3H₂ ⇌ 2NH₃)
- Nickel (Ni): catalyst in the hydrogenation of vegetable oils to make margarine (C=C + H₂ → C–C)
- Vanadium(V) oxide (V₂O₅): catalyst in the Contact process for sulfuric acid manufacture
- Platinum (Pt) and palladium (Pd): catalysts in catalytic converters, converting CO and NO₂ to CO₂ and N₂
What are the everyday uses of transition metals?
The combination of high strength, high melting points, electrical conductivity, and chemical stability makes transition metals indispensable in manufacturing:
- Iron/steel: construction, vehicles, machinery (by far the most widely used metal)
- Copper: electrical wiring, plumbing, coins (excellent conductor, resistant to corrosion)
- Titanium: aerospace and medical implants (strong, low density, biocompatible)
- Chromium: stainless steel alloy, chrome plating (prevents corrosion)
- Tungsten: light bulb filaments, cutting tools (highest melting point of any metal: 3,422 °C)
- Nickel: rechargeable batteries, stainless steel alloy
Frequently asked questions
What are the transition metals in GCSE chemistry?
The transition metals are the d-block elements in the central section of the periodic table, including iron (Fe), copper (Cu), nickel (Ni), chromium (Cr), manganese (Mn), cobalt (Co), titanium (Ti), and zinc (Zn). They are characterised by high melting points and densities, variable oxidation states, the ability to form coloured compounds, and catalytic activity. They contrast sharply with Group 1 alkali metals, which are soft, reactive, and low-melting.
Why are transition metal compounds coloured?
Transition metal ions have partially filled d-orbitals. When visible light falls on a solution or crystal of a transition metal compound, electrons absorb specific wavelengths to jump between d-orbital energy levels. The wavelengths that are not absorbed are reflected or transmitted — those are the colours we see. Because the energy gap between d-orbitals varies with the metal ion and the surrounding ligands, different transition metal compounds absorb different colours and therefore appear different colours.
What does variable oxidation state mean?
An oxidation state describes how many electrons an atom has lost (positive values) or gained (negative values) in forming a compound. Transition metals can lose different numbers of electrons depending on conditions, giving them two or more stable oxidation states. Iron, for instance, can be Fe²⁺ (iron(II)) or Fe³⁺ (iron(III)). The oxidation state is shown by the Roman numeral in the compound's name. Variable oxidation states are possible because the d-electrons and outermost s-electrons are close in energy and can be removed selectively.
How is iron used as a catalyst in the Haber process?
In the Haber process, nitrogen (N₂) and hydrogen (H₂) are passed over an iron catalyst at about 450 °C and 200 atmospheres pressure to produce ammonia (N₂ + 3H₂ ⇌ 2NH₃). The iron catalyst provides a surface on which the N≡N triple bond — very stable and difficult to break — adsorbs and weakens, making the reaction feasible at an industrially practical temperature. Without the catalyst, the reaction would be far too slow at a temperature that favours ammonia formation. Iron's variable oxidation states and d-electron availability are what make it effective in this role.
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