Last updated: 14 October 2026
The first-row transition elements occupy the d-block of the Periodic Table, but being in the d-block does not automatically make an element a transition element. The crucial question is whether an incomplete d subshell is present in the atom itself or in one of its cations.
Once that distinction is clear, many of the familiar properties of transition elements begin to fit together. Their electronic configurations explain why their atomic sizes change relatively little across the series, why many of them form ions in several oxidation states, and why transition metals and their compounds are so useful as catalysts.
Try these five questions before reading the chapter. Pay particular attention to electronic configurations and the distinction between d-block elements and transition elements.
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What Exactly Is a Transition Element?
Transition elements are found in the d-block of the Periodic Table, but the terms d-block element and transition element are not interchangeable.
For the current 9476 syllabus, the definition deliberately considers both the electronic configuration of the atom and the electronic configurations of its cations.
A transition element is a d-block element whose atom has an incomplete d subshell, or which can give rise to cations with an incomplete d subshell.
Why scandium now counts as a transition element
A neutral scandium atom has the electronic configuration:
Its 3d subshell contains only one electron, so the d subshell of the atom is incomplete. Scandium therefore satisfies the current 9476 definition.
Its common ion, Sc³⁺, has the configuration:
Sc³⁺ has an empty d subshell, but that does not matter for the current definition because the Sc atom itself already satisfies the requirement.
What changed from the previous 9729 syllabus?
Previous 9729 approach: the definition focused on a d-block element forming one or more stable ions with an incompletely filled d subshell. Scandium was therefore excluded because its stable Sc³⁺ ion is 3d⁰. The transition elements normally considered ran from Ti to Cu.
Current 9476 approach: the atom itself may satisfy the definition. Since Sc is [Ar] 3d¹4s², scandium is now included, and the syllabus refers to the first set of transition elements from Sc to Cu.
Exam tip: If you learnt the older definition from previous notes or older examination materials, do not use it for 9476. For the current syllabus, remember the two possibilities: incomplete d subshell in the atom OR in a cation.
Why zinc is still not a transition element
Zinc provides the useful contrast. Its atom and its common ion have the following configurations:
Zn²⁺: [Ar] 3d¹⁰
The 3d subshell is completely filled in both Zn and Zn²⁺. Zinc is therefore a d-block element but not a transition element.
Why copper is a transition element
A copper atom has a complete 3d subshell:
However, Cu²⁺ has the configuration:
Cu²⁺ therefore contains an incomplete d subshell, so copper satisfies the transition-element definition.
| Element | Relevant configuration | Transition element under 9476? |
|---|---|---|
| Sc | Sc atom: [Ar] 3d¹4s² | Yes — the atom has an incomplete d subshell |
| Cu | Cu²⁺: [Ar] 3d⁹ | Yes — a cation has an incomplete d subshell |
| Zn | Zn: [Ar] 3d¹⁰4s²; Zn²⁺: [Ar] 3d¹⁰ | No — neither has an incomplete d subshell |
Electronic Configurations of the First-Row Transition Elements
Across the first transition series, electrons are added mainly to the 3d subshell. The simple expected pattern is therefore based on:
Two important exceptions are chromium and copper. Their experimentally observed configurations contain only one 4s electron.
| Element | Atomic number | Electronic configuration |
|---|---|---|
| Sc | 21 | [Ar] 3d¹4s² |
| Ti | 22 | [Ar] 3d²4s² |
| V | 23 | [Ar] 3d³4s² |
| Cr | 24 | [Ar] 3d⁵4s¹ |
| Mn | 25 | [Ar] 3d⁵4s² |
| Fe | 26 | [Ar] 3d⁶4s² |
| Co | 27 | [Ar] 3d⁷4s² |
| Ni | 28 | [Ar] 3d⁸4s² |
| Cu | 29 | [Ar] 3d¹⁰4s¹ |
The chromium and copper exceptions
If the simple filling pattern were followed mechanically, chromium might be expected to be [Ar] 3d⁴4s² and copper [Ar] 3d⁹4s². Their actual configurations are:
Cu: [Ar] 3d¹⁰4s¹
At A-Level, you should know these observed configurations. The energies of the 3d and 4s orbitals are very close, so the simple filling-order rule should not be treated as an absolute law once these orbitals are occupied.
Do not predict chromium and copper by blindly applying the Aufbau pattern. Memorise their observed configurations: Cr = [Ar] 3d⁵4s¹ and Cu = [Ar] 3d¹⁰4s¹.
Forming cations: remove 4s electrons first
There is another rule that often surprises students. Although the 4s orbital is filled before 3d when constructing the electronic configurations of neutral atoms, the 4s electrons are removed first when transition-metal cations are formed.
Fe is [Ar] 3d⁶4s².
Fe²⁺ is formed by removing the two 4s electrons:
Fe²⁺ = [Ar] 3d⁶
Removing one additional 3d electron gives:
Fe³⁺ = [Ar] 3d⁵
Similarly:
Cu⁺: [Ar] 3d¹⁰
Cu²⁺: [Ar] 3d⁹
Why Atomic Radii Are Relatively Invariant
Moving from Sc to Cu, the proton number increases. If this were the only change, the increasing nuclear charge would pull the electrons progressively closer to the nucleus and atomic radius would decrease substantially.
However, the additional electrons are being added mainly to the 3d subshell. These additional 3d electrons provide greater shielding from the increasing nuclear charge.
Across the transition series, increasing nuclear charge tends to decrease atomic radius, while increasing shielding by additional 3d electrons opposes this effect. The atomic radii are therefore relatively invariant.
The cancellation is not exact, so the atomic radii are not literally identical. The important syllabus point is that their variation is much smaller than might otherwise be expected as proton number increases.
Why First Ionisation Energies Are Relatively Invariant
The same competition helps explain why the first ionisation energies of the transition elements also show comparatively little overall change across the series.
Increasing nuclear charge: increases the attraction between the nucleus and the electron being removed, tending to increase first ionisation energy.
Increasing shielding: additional 3d electrons shield the outer electron from some of the increasing nuclear charge, reducing the size of that increase.
Consequently, first ionisation energies fluctuate across the series rather than increasing dramatically and regularly. They are described as relatively invariant, not exactly constant.
Melting Points and Densities Compared with Calcium
Calcium is a useful representative s-block metal for comparison. In general, the first-row transition elements have considerably higher melting points and densities than calcium.
| Property | Transition elements | Comparison with calcium |
|---|---|---|
| Melting point | Generally high | Generally much higher |
| Density | Generally high | Much higher |
| Metallic bonding | Strong | Generally stronger |
Why are their melting points generally higher?
Transition metals can involve both 4s and 3d electrons in metallic bonding. This generally produces stronger attractions between the positive metal centres and the delocalised electrons than in calcium.
More energy is therefore required to overcome the metallic bonding during melting, giving transition metals generally higher melting points.
Why are their densities higher?
Across the transition series, relative atomic mass increases while atomic radii remain comparatively small. A relatively large mass is therefore packed into a relatively small volume.
Higher density arises from the combination of relatively high atomic masses and relatively small atomic sizes.
Why Transition Elements Have Variable Oxidation States
One of the most characteristic chemical properties of transition elements is their ability to form compounds in more than one oxidation state. Iron, for example, commonly forms Fe²⁺ and Fe³⁺, while copper commonly forms Cu⁺ and Cu²⁺.
This happens because the 3d and 4s subshells are relatively close in energy. The 4s electrons and different numbers of 3d electrons can therefore participate in bonding or be removed, allowing the same element to exist in several oxidation states.
Variable oxidation states arise because the 3d and 4s subshells have similar energies, so different numbers of 3d and 4s electrons can participate in bonding or be removed.
Predicting possible oxidation states from electronic configuration
A useful A-Level guideline for predicting the highest oxidation state of a first-row transition element is to count the 4s electrons together with the unpaired 3d electrons.
The word unpaired matters. From Sc to Mn, electrons enter separate 3d orbitals before pairing, so the number of unpaired 3d electrons increases. After Mn, additional 3d electrons must begin pairing, so the number of unpaired 3d electrons decreases.
This gives the familiar rise in the predicted maximum oxidation state from Sc to Mn, followed by a fall from Fe onwards.
Vanadium: [Ar] 3d³4s²
The three 3d electrons are unpaired.
3 unpaired 3d electrons + 2 electrons in 4s = 5
Therefore, vanadium can reach an oxidation state of +5.
Manganese: [Ar] 3d⁵4s²
All five 3d electrons are unpaired.
5 unpaired 3d electrons + 2 electrons in 4s = 7
Therefore, manganese can reach an oxidation state of +7, as in MnO₄⁻.
Why the maximum oxidation state falls after manganese
A 3d subshell contains five orbitals. From Sc to Mn, electrons occupy these orbitals singly, so the number of unpaired 3d electrons rises from one to five.
Beginning with Fe, the 3d electrons start to pair. The number of unpaired 3d electrons therefore decreases:
Co: 3d⁷ → 3 unpaired electrons
Ni: 3d⁸ → 2 unpaired electrons
Applying the guideline gives predicted highest oxidation states of +6 for Fe, +5 for Co and +4 for Ni.
| Element | Electronic configuration | Unpaired 3d electrons | 4s electrons | Predicted highest oxidation state |
|---|---|---|---|---|
| Sc | [Ar] 3d¹4s² | 1 | 2 | +3 |
| Ti | [Ar] 3d²4s² | 2 | 2 | +4 |
| V | [Ar] 3d³4s² | 3 | 2 | +5 |
| Cr | [Ar] 3d⁵4s¹ | 5 | 1 | +6 |
| Mn | [Ar] 3d⁵4s² | 5 | 2 | +7 |
| Fe | [Ar] 3d⁶4s² | 4 | 2 | +6 |
| Co | [Ar] 3d⁷4s² | 3 | 2 | +5 |
| Ni | [Ar] 3d⁸4s² | 2 | 2 | +4 |
| Cu | [Ar] 3d¹⁰4s¹ | 0 | 1 | +1 by the simple guideline; Cu is an exception |
Do not confuse the highest oxidation state with the most common oxidation state
The guideline predicts how high an oxidation state may be possible from the electronic configuration. It does not mean that this oxidation state is necessarily the most stable or most commonly encountered.
For example, the guideline allows Fe to reach +6, but Fe²⁺ and Fe³⁺ are much more commonly encountered in A-Level Chemistry. Similarly, manganese can reach +7 in MnO₄⁻, but Mn²⁺ is also an important and stable oxidation state.
Predicted highest oxidation state: use the electronic configuration and count the 4s electrons plus unpaired 3d electrons as a useful guide.
Common oxidation states: depend on the element and its chemical environment. A lower oxidation state may be much more stable or commonly encountered.
Copper is the important exception
Copper shows why the guideline must not be treated as an absolute law. Its electronic configuration is:
All ten 3d electrons are paired, so copper has zero unpaired 3d electrons and one 4s electron. The simple guideline would therefore suggest +1:
Copper does indeed form Cu⁺, but it also commonly forms Cu²⁺. Formation of Cu²⁺ requires removal of one additional electron from the filled 3d subshell:
Cu²⁺: [Ar] 3d⁹
Treat “4s electrons + unpaired 3d electrons” as a useful prediction guideline, not an absolute law. Copper is the important exception: although the simple count gives +1, Cu²⁺ is also a common oxidation state.
Why Transition Metals and Their Compounds Are Good Catalysts
A catalyst provides an alternative reaction pathway with a lower activation energy and is regenerated at the end of the reaction. Transition elements can be particularly effective catalysts for two related but distinct reasons.
Homogeneous catalysis: changing oxidation states
In homogeneous catalysis, the catalyst and reactants are in the same phase. A transition-metal ion can often change readily between different oxidation states.
It can therefore participate in one electron-transfer step, form an intermediate oxidation state, and then take part in another step that regenerates the original catalyst.
Variable oxidation states allow transition-metal ions to provide an alternative sequence of redox steps with a lower activation energy.
Heterogeneous catalysis: adsorption onto a metal surface
In heterogeneous catalysis, the catalyst is in a different phase from the reactants. A solid transition metal provides a surface onto which reactant particles can adsorb.
Adsorption brings reactant particles close together and can weaken bonds within the adsorbed molecules. Reaction can then occur more readily on the surface before the products desorb.
Example: iron is used as the heterogeneous catalyst in the Haber process. Nitrogen and hydrogen adsorb onto the iron surface, where their bonds are weakened and reaction can occur by a lower-energy pathway.
Homogeneous transition-metal catalyst: variable oxidation states can allow successive electron-transfer steps.
Heterogeneous transition-metal catalyst: reactants adsorb onto the catalyst surface, where bonds may be weakened and reacting particles are brought together.
Common Mistakes
| Common mistake | Why it is wrong | Better approach |
|---|---|---|
| Defining a transition element only by whether it forms a cation with an incomplete d subshell | That omits the first part of the current 9476 definition and would incorrectly exclude scandium. | State that either the atom itself or a cation may possess an incomplete d subshell. |
| Saying every d-block element is a transition element | Zn is in the d-block but Zn and Zn²⁺ both have a complete 3d¹⁰ subshell. | Test the actual electronic configurations against the definition. |
| Removing 3d electrons before 4s electrons when forming cations | For transition-metal cations, 4s electrons are removed before 3d electrons. | Write the atomic configuration first, remove 4s electrons, then remove 3d electrons if necessary. |
| Adding all the 3d electrons when predicting the highest oxidation state | The simple guideline uses the number of unpaired 3d electrons, not the total number of 3d electrons. | Count the 4s electrons and the unpaired 3d electrons, then remember that this is only a prediction guideline. |
| Assuming the predicted highest oxidation state must be the most common | For example, Fe can reach +6, but Fe²⁺ and Fe³⁺ are much more commonly encountered. | Distinguish between a possible maximum oxidation state and the oxidation states most stable in a particular chemical environment. |
| Applying the oxidation-state guideline rigidly to Cu | Cu is [Ar] 3d¹⁰4s¹, so the simple count gives +1, yet Cu²⁺ is also a common oxidation state. | Remember Cu as an important exception and know both Cu⁺ and Cu²⁺. |
What you should be able to do
Define a transition element
Use the current 9476 definition involving an incomplete d subshell in the atom or in a cation.
Write electronic configurations
State the configurations of first-row transition atoms and ions, including the Cr and Cu exceptions.
Form transition-metal cations
Remove 4s electrons before 3d electrons when deriving ionic configurations.
Explain periodic trends
Explain why atomic radii and first ionisation energies are relatively invariant across the transition series.
Predict oxidation states
Use the number of 4s electrons plus unpaired 3d electrons as a useful guide to the highest oxidation state, while recognising its limitations and the Cu exception.
Explain catalytic behaviour
Distinguish between catalysis involving variable oxidation states and heterogeneous catalysis involving adsorption onto a surface.
These ten questions test whether you can apply the definition, electronic configurations, periodic trends, oxidation states and catalytic behaviour of transition elements.
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Frequently asked questions
Questions Students Often Ask About Transition Elements
The definition change involving scandium is likely to be the first source of confusion, but several other questions about electronic configurations and oxidation states naturally follow from it.
Is scandium really a transition element in the 9476 syllabus?
Yes. Sc has the atomic configuration [Ar] 3d¹4s², so the Sc atom itself has an incomplete d subshell. The current definition allows either the atom or one of its cations to possess an incomplete d subshell. Sc therefore qualifies even though Sc³⁺ is [Ar] 3d⁰.
Why is zinc not a transition element if zinc is in the d-block?
Zn is [Ar] 3d¹⁰4s² and Zn²⁺ is [Ar] 3d¹⁰. The d subshell is completely filled in both the atom and its common cation, so zinc does not satisfy the transition-element definition. This is why d-block element and transition element do not mean exactly the same thing.
Why do we fill 4s before 3d but remove 4s electrons first?
The relative energies of the 4s and 3d orbitals change as electrons are added to the atom. Although 4s is occupied before 3d when the neutral atoms are built up, once the 3d subshell is occupied the 4s electrons are removed first during cation formation. For examination purposes, apply the rule directly: remove 4s electrons before 3d electrons.
Is the maximum oxidation state always the number of 4s electrons plus unpaired 3d electrons?
It is a useful A-Level prediction guideline, but it should not be treated as an absolute law. It works particularly well for seeing the rise in maximum oxidation state from Sc to Mn and the subsequent fall as 3d electrons begin to pair. Copper is an important exception: Cu is [Ar] 3d¹⁰4s¹, so the simple count predicts +1, yet Cu²⁺ is also a common oxidation state.
Does a high predicted oxidation state mean that it is the most common oxidation state?
No. The guideline tells you about a possible high oxidation state, not necessarily the most stable or most commonly encountered one. For example, Fe can reach +6, but Fe²⁺ and Fe³⁺ are far more commonly encountered in A-Level Chemistry. The stability of an oxidation state depends on the chemical environment.
The key to this chapter is the d subshell. It explains why Sc qualifies under the current definition, why Zn does not, why transition elements can form several oxidation states, and why many of them are useful catalysts. When predicting oxidation states, electronic configuration gives us a useful starting point, but the chemical environment ultimately determines which oxidation states are stable. With those ideas established, the next step is to see how d electrons also explain one of the most recognisable features of transition chemistry: the colours of transition-metal complexes.