Electronic Configuration of Copper (Cu) — Exception to Aufbau Principle

Atomic number of Copper (Cu): 29 Total electrons: 29

Full electronic configuration: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹

Noble gas (abbreviated) configuration: [Ar] 3d¹⁰ 4s¹ (where [Ar] = 1s² 2s² 2p⁶ 3s² 3p⁶, representing the Argon core with 18 electrons)

Subshell notation: 29 electrons distributed as:

• n=1: 1s² (2 electrons)
• n=2: 2s²2p⁶ (8 electrons)
• n=3: 3s²3p⁶3d¹⁰ (18 electrons)
• n=4: 4s¹ (1 electron) Total: 2 + 8 + 18 + 1 = 29 ✓

The Aufbau principle states that electrons fill orbitals in order of increasing energy. For transition metals, 4s fills before 3d. Based on this:

Expected configuration (by Aufbau): [Ar] 3d⁹ 4s² Actual configuration: [Ar] 3d¹⁰ 4s¹

Reason for the exception:

1. Extra stability of fully filled d-orbitals: A completely filled 3d¹⁰ configuration is energetically more stable than a partially filled 3d⁹ configuration. This is due to:
   • Symmetrical charge distribution (all five d-orbitals fully occupied)
   • Maximum exchange energy (electrons in same subshell with same spin can exchange — more pairs = more exchange energy = lower energy state)
2. The energy difference between 3d and 4s is small in Cu, making the 3d¹⁰4s¹ arrangement more stable than 3d⁹4s².

Chromium (Cr, Z=24) has a similar exception: [Ar] 3d⁵ 4s¹ (half-filled d-orbital stability).

Copper forms two common ions by losing electrons from the outer shell:

Cu⁺ (copper(I) ion): loses the 4s¹ electron first Configuration: [Ar] 3d¹⁰ (completely stable, closed d-shell)

Cu²⁺ (copper(II) ion): loses 4s¹ and one 3d electron Configuration: [Ar] 3d⁹

Note: In ionic species, 4s electrons are removed before 3d electrons, because once nuclear charge exceeds the electron count, 3d is stabilised below 4s.

Cu²⁺ is the more common ion in aqueous solution (e.g., CuSO₄ solution is blue due to [Cu(H₂O)₆]²⁺). Cu⁺ is less stable in aqueous solution and disproportionates: 2Cu⁺ → Cu + Cu²⁺

The orbital box diagram for copper [Ar] 3d¹⁰ 4s¹:

3d orbitals (5 boxes, all fully filled): [↑↓] [↑↓] [↑↓] [↑↓] [↑↓] ← 3d¹⁰ (10 electrons, all paired)

4s orbital (1 box, one electron): [↑ ] ← 4s¹ (1 unpaired electron)

Key observations:

• All 5 d-orbitals are fully occupied → extra stability
• One unpaired electron in 4s → Cu is paramagnetic (weakly attracted to magnetic fields)
• Cu²⁺ has one unpaired electron in 3d⁹ → also paramagnetic
• Copper's characteristic reddish colour is due to electron transitions involving the 3d¹⁰ electrons

1. Multiple oxidation states: Cu shows +1 and +2 oxidation states due to the ease of losing 4s and 3d electrons.

2. Coloured compounds: Cu²⁺ compounds are typically blue/green because the partially filled 3d⁹ subshell undergoes d-d electronic transitions absorbing visible light.

3. Catalytic activity: The partially filled d-orbitals allow Cu to act as a catalyst.

4. Good electrical conductivity: The single 4s¹ electron is loosely held and delocalised in metallic bonding, making Cu an excellent electrical and thermal conductor.

5. Complex formation: Cu²⁺ readily forms coordination complexes (e.g., [Cu(NH₃)₄]²⁺ — deep blue Schweizer's reagent).

6. Diamagnetic Cu⁺: Has 3d¹⁰ configuration — all electrons paired → diamagnetic.