Transition Metal Electron Configurations: 4s, 3d and Ions

The first row of transition metals, from scandium (21) to zinc (30), is where electron configurations stop being routine. The 3d and 4s subshells are so close in energy that small effects decide which fills first, two elements break the expected pattern, and ions lose electrons in a different order from the one in which atoms gained them. This article works through the whole row element by element, then shows how the configurations explain the chemistry that makes transition metals special.

Background: the Aufbau principle, Hund’s rule and noble gas shorthand.

The whole first row

After argon (18 electrons), potassium and calcium fill 4s. From scandium onward, electrons enter the 3d subshell.

Element	Z	Configuration	Unpaired electrons
Scandium, Sc	21	[Ar] 3d¹ 4s²	1
Titanium, Ti	22	[Ar] 3d² 4s²	2
Vanadium, V	23	[Ar] 3d³ 4s²	3
Chromium, Cr	24	[Ar] 3d⁵ 4s¹	6
Manganese, Mn	25	[Ar] 3d⁵ 4s²	5
Iron, Fe	26	[Ar] 3d⁶ 4s²	4
Cobalt, Co	27	[Ar] 3d⁷ 4s²	3
Nickel, Ni	28	[Ar] 3d⁸ 4s²	2
Copper, Cu	29	[Ar] 3d¹⁰ 4s¹	1
Zinc, Zn	30	[Ar] 3d¹⁰ 4s²	0

Configurations are written in shell order (3d before 4s), which matches the order in which electrons are removed when ions form.

Worked example 1: vanadium

1. Z = 23, so 23 electrons.
2. [Ar] accounts for 18; 5 remain.
3. Fill 4s² (2), then 3d³ (3).
4. V: [Ar] 3d³ 4s²

By Hund’s rule, the three 3d electrons occupy separate orbitals with parallel spins: 3 unpaired.

Worked example 2: the chromium exception

The simple filling order predicts [Ar] 3d⁴ 4s². The actual ground state is [Ar] 3d⁵ 4s¹.

Why? The 3d and 4s orbitals are very close in energy in chromium. Moving one electron from 4s into 3d gives a half-filled d subshell, five electrons with parallel spins in five separate orbitals. This arrangement:

• reduces electron–electron repulsion (the 4s electron no longer shares an orbital), and
• maximises the stabilising exchange energy between parallel spins.

The small energy cost of the promotion is outweighed, so 3d⁵ 4s¹ is lower in energy. See electron configuration exceptions.

Worked example 3: the copper exception

Predicted: [Ar] 3d⁹ 4s². Actual: [Ar] 3d¹⁰ 4s¹.

A completely filled d subshell is especially stable; moving one electron from 4s to 3d achieves it. The effect is related to the fact that, as nuclear charge increases across the row, the 3d orbitals drop in energy more steeply than 4s.

Why 3d and 4s change places

In potassium and calcium, 4s is lower than 3d, because 4s electrons penetrate close to the nucleus while 3d electrons are well shielded. See effective nuclear charge.

But as protons are added across the transition series, the 3d orbitals fall in energy faster than 4s, because d electrons are held increasingly tightly by the growing nuclear charge. Once 3d is occupied, it lies below 4s. That’s the key to the next section.

Writing ions: 4s electrons go first

When transition metals form positive ions, electrons are removed from the 4s subshell first, then from 3d.

Worked example 4: iron ions

• Fe: [Ar] 3d⁶ 4s²
• Fe²⁺: remove 2 from 4s → [Ar] 3d⁶ (4 unpaired)
• Fe³⁺: remove 1 more from 3d → [Ar] 3d⁵ (5 unpaired, half-filled)

Worked example 5: chromium(III)

• Cr: [Ar] 3d⁵ 4s¹
• Cr³⁺: remove 1 from 4s and 2 from 3d → [Ar] 3d³

Worked example 6: copper ions

• Cu: [Ar] 3d¹⁰ 4s¹
• Cu⁺: remove the 4s electron → [Ar] 3d¹⁰
• Cu²⁺: remove one 3d electron → [Ar] 3d⁹

Worked example 7: scandium and zinc

• Sc³⁺: [Ar] (no d electrons)
• Zn²⁺: [Ar] 3d¹⁰ (full d subshell)

Because scandium’s only common ion has no d electrons and zinc’s has a full d subshell, neither forms an ion with a partly filled d subshell. Under the usual IUPAC-style definition used in many courses, that’s why zinc (and sometimes scandium) isn’t counted as a true transition metal, even though both sit in the d-block. See transition metals.

How configurations explain transition metal chemistry

Variable oxidation states

The 4s and 3d electrons are close in energy, so different numbers can be removed or shared with only gradual increases in energy. Manganese, for example, shows oxidation states from +2 (Mn²⁺, [Ar] 3d⁵) up to +7 (in the permanganate ion, MnO₄⁻, where all seven 4s and 3d electrons are involved in bonding). See how to find oxidation numbers.

Element	Common oxidation states
Ti	+2, +3, +4
V	+2, +3, +4, +5
Cr	+2, +3, +6
Mn	+2, +4, +7
Fe	+2, +3
Co	+2, +3
Ni	+2
Cu	+1, +2
Zn	+2

Colour

In complexes, surrounding ligands split the five d orbitals into groups with slightly different energies. Electrons can absorb visible light and jump between them. Ions with partly filled d subshells are usually coloured:

• Cu²⁺ (3d⁹): blue in water
• Ni²⁺ (3d⁸): green
• Co²⁺ (3d⁷): pink
• Fe³⁺ (3d⁵): yellow to orange-brown

Ions with empty (Sc³⁺, Ti⁴⁺) or full (Zn²⁺, Cu⁺) d subshells have no such transitions and are usually colourless or white.

Magnetism

Unpaired d electrons make ions paramagnetic. Fe³⁺ and Mn²⁺, with five unpaired electrons each, are strongly paramagnetic. See paramagnetic vs diamagnetic and how to count unpaired electrons.

Catalysis

The ability to switch between oxidation states, and to use d orbitals to bind other molecules, makes many transition metals excellent catalysts: iron in the Haber process, vanadium(V) oxide in the Contact process for sulfuric acid, nickel in hydrogenating vegetable oils. See reaction rates and catalysts.

Beyond the first row

In the second and third transition rows, the energies of the d and s subshells are even closer, and more “irregular” configurations appear:

• Molybdenum: [Kr] 4d⁵ 5s¹ (like chromium)
• Silver: [Kr] 4d¹⁰ 5s¹ and gold: [Xe] 4f¹⁴ 5d¹⁰ 6s¹ (like copper)
• Palladium: [Kr] 4d¹⁰, with no 5s electrons at all
• Niobium: [Kr] 4d⁴ 5s¹; ruthenium: [Kr] 4d⁷ 5s¹

For heavy elements, relativistic effects also shift orbital energies. See relativistic effects in chemistry. Always check heavier configurations in a reliable data source, such as each element’s page, e.g. palladium.

Practice

Write the configurations of:

1. Ti²⁺
2. Mn³⁺
3. Co³⁺
4. Ni²⁺
5. V⁵⁺

Answers: (1) [Ar] 3d²; (2) [Ar] 3d⁴; (3) [Ar] 3d⁶; (4) [Ar] 3d⁸; (5) [Ar].

A quick way to check any answer: the number of d electrons in a first-row transition metal ion equals the group number minus the charge. Fe³⁺ (group 8): 8 − 3 = 5 d electrons ✓. Cu²⁺ (group 11): 11 − 2 = 9 ✓.

Key takeaways

• From Sc to Zn, electrons fill 3d after 4s; configurations run from [Ar] 3d¹ 4s² to [Ar] 3d¹⁰ 4s².
• Chromium (3d⁵ 4s¹) and copper (3d¹⁰ 4s¹) are exceptions, favouring half-filled and filled d subshells.
• Once 3d is occupied, it lies below 4s, so ions lose 4s electrons first.
• Close d and s energies give variable oxidation states; partly filled d subshells give colour and paramagnetism.
• Heavier transition metals show even more irregular configurations.