Transition metals: the d-block elements and their chemistry
Explore the d-block: variable oxidation states, colored ions, magnetism, catalysis, and why transition metals form so many complexes.
Quick answer
Transition metals are the d-block elements whose atoms or common ions have partially filled d subshells, giving them variable oxidation states, colored compounds, paramagnetism, and rich coordination chemistry.
Example: Iron commonly shows +2 (pale green Fe2+) and +3 (yellow-brown Fe3+), forms complexes like [Fe(CN)6]4-, and catalyzes the Haber process: classic transition-metal behavior.
What you will learn
- Define a transition metal using the incomplete-d-subshell criterion
- Write electron configurations for first-row transition metals including the Cr and Cu exceptions
- Explain the origin of variable oxidation states
- Relate color and paramagnetism to d-electron count
- Describe why transition metals form complexes and act as catalysts
- Summarize key trends across a transition series
Which elements count as transition metals
The transition metals occupy the d-block of the periodic table, groups 3 through 12. But the precise definition is narrower than just location: a transition metal is an element whose atom has a partially filled d subshell, or which can form cations with an incomplete d subshell.
This strict definition excludes zinc, cadmium, and mercury. Zinc's atom is [Ar]3d10 4s2 and its only common ion, Zn2+, is also 3d10, so its d subshell is never incomplete. Scandium qualifies through its atom ([Ar]3d1 4s2) even though Sc3+ is d0. In casual speech chemists often say transition metal for any d-block element, but the strict definition is what exams expect.
An element counts as a transition metal only if its atom or at least one common ion has an incomplete d subshell. Zn, Cd, and Hg fail this test.
Electron configurations and the famous exceptions
Filling order puts the ns orbital before (n-1)d, giving configurations like Sc [Ar]4s2 3d1 through Zn [Ar]4s2 3d10. Two famous exceptions break the pattern: chromium is [Ar]4s1 3d5 and copper is [Ar]4s1 3d10.
Each exception trades a 4s electron for extra d stability: a half-filled d5 subshell for chromium and a filled d10 subshell for copper gain enough exchange stabilization to outweigh the promotion cost. These are the only such anomalies in the first transition series worth memorizing.
When transition metals ionize, the ns electrons leave first. Fe becomes Fe2+ ([Ar]3d6) by losing its two 4s electrons, because the 3d orbitals sink below 4s in energy once electrons are removed.
Variable oxidation states
Because the ns and (n-1)d electrons are close in energy, transition metals can lose different numbers of electrons and access many oxidation states. Iron shows +2 and +3; chromium shows +2, +3, and +6; vanadium runs from +2 to +5; manganese spans +2 all the way to +7 in the deep-purple permanganate ion, MnO4-.
Higher oxidation states become more stable toward the right of the early transition series and are strongly oxidizing: dichromate (Cr2O7 2-) and permanganate are powerful oxidizers precisely because the metal is electron-hungry. Lower states dominate for the later metals, where removing d electrons costs more energy.
Color and magnetism
Most transition-metal ions are colored because their partially filled d subshells allow d-d electron transitions that absorb visible light. Aqueous Cu2+ is blue, Co2+ pink, Ni2+ green, and MnO4- an intense purple, though that last color comes from charge transfer rather than a d-d jump.
Unpaired d electrons also make these ions paramagnetic. High-spin Mn2+ and Fe3+, both d5, have five unpaired electrons and are strongly attracted to magnets, while d0 ions like Sc3+ and d10 ions like Zn2+ are colorless and diamagnetic.
Complex formation and catalysis
Transition metals form complexes readily because their ions are small, often highly charged, and offer vacant d orbitals for ligand electron pairs. Nearly every transition metal appears in coordination compounds, from [Ag(NH3)2]+ to [Fe(CN)6]4-.
The same properties make them outstanding catalysts. They adsorb reactants onto their surfaces, cycle between oxidation states to transfer electrons, and stabilize transition states through complex formation. Industrial examples include iron in the Haber synthesis of ammonia, nickel in the hydrogenation of vegetable oils, vanadium(V) oxide in the Contact process for sulfuric acid, and platinum-group metals in catalytic converters.
Trends across the d-block
Moving left to right across a transition series, atomic radius contracts as rising nuclear charge pulls the diffuse d electrons inward, though the contraction levels off and even reverses slightly near the end. Ionization energies climb gradually, densities and melting points peak near the middle of the series where metallic bonding is strongest.
Reactivity falls across the series: early metals like titanium and vanadium are reactive, middle metals like iron and cobalt moderately so, and late metals like gold and platinum are noble enough to occur native. These trends explain why the chemistry of the d-block feels like a gradual transition from reactive metals to near-noble ones.
Transition metals in everyday life
Transition metals are woven into modern life. Iron and steel frame buildings and vehicles, chromium makes stainless steel stainless, and copper wiring carries electricity everywhere.
Cobalt, nickel, and manganese oxides store energy in lithium-ion batteries, titanium dioxide whitens paint and sunscreen, and silver compounds once dominated photography. From pigments to catalysts to electronics, the d-block supplies the materials of the industrial world.
Variables and units
| Symbol | Name | Unit | Meaning |
|---|---|---|---|
| d-electron count | dimensionless | Number of electrons in the (n-1)d subshell, which determines color, magnetism, and spin state |
Common mistakes
The strict definition requires an incomplete d subshell in the atom or a common ion. Zinc, cadmium, and mercury have full d10 configurations in every common state and are excluded, even though they sit in the d-block.
Color needs a partially filled d subshell for d-d transitions. Sc3+ and Ti4+ are d0, while Zn2+ and Cu+ are d10, and all are colorless.
The ns electrons are lost first because the (n-1)d orbitals drop below ns in energy once the atom is ionized. Fe2+ is [Ar]3d6, formed by losing the two 4s electrons, not d electrons.
Mercury is a liquid at room temperature, and metals like zinc and cadmium are comparatively soft with low melting points. The trend toward hard, high-melting metals is strongest in the middle of each series.
Real-world applications
Iron and carbon make steel, and added chromium makes stainless steel resist corrosion
Iron catalysts drive the Haber process that produces ammonia for fertilizer
Lithium-ion battery cathodes rely on cobalt, nickel, and manganese oxides
Platinum, palladium, and rhodium in catalytic converters clean up vehicle exhaust
Titanium dioxide is the brilliant white pigment in paints, paper, and sunscreen
Frequently asked questions
Why do chromium and copper have unusual electron configurations?
Chromium is [Ar]4s1 3d5 rather than [Ar]4s2 3d4 because a half-filled d subshell gains extra exchange stabilization that outweighs the cost of promoting one 4s electron. Copper is [Ar]4s1 3d10 for the analogous reason: a completely filled d subshell is especially stable. These are the two famous exceptions in the first transition series.
Why do transition metals show so many oxidation states?
Transition metals can lose their ns electrons and then a variable number of (n-1)d electrons, since the two subshells are close in energy. Iron loses its 4s electrons to give Fe2+ and one more 3d electron to give Fe3+; manganese can lose up to seven electrons to reach the +7 state in permanganate.
Why is zinc not considered a transition metal?
By the strict definition, a transition metal must have an incomplete d subshell in its atom or in at least one common ion. Zinc is [Ar]3d10 4s2 and forms only Zn2+, which is also 3d10, so its d subshell is never incomplete. Cadmium and mercury are excluded for the same reason, though all three sit in the d-block.
Why are so many transition-metal compounds colored?
Color in these ions comes from d-d electron transitions: visible light promotes an electron between split d orbitals. Ions with partially filled d subshells, like Cu2+ or Co2+, absorb visible wavelengths. Ions with d0 (Sc3+, Ti4+) or d10 (Zn2+, Cu+) configurations have no possible d-d transition and are colorless.
Why are transition metals such good catalysts?
Good catalysts need to bind reactants, shuffle electrons, and release products, and transition metals do all three. They form complexes easily, cycle between oxidation states to carry electrons, and provide surfaces where bonds break and reform. Iron in the Haber process, nickel in hydrogenation, and platinum in catalytic converters all work this way.