Copper (II) ions are able to participate in co-ordinate covalent bonding because they

Analyze the question: It asks why Cu²⁺ ions can form coordinate (dative) covalent bonds, for example in [Cu(NH₃)₄]²⁺ or [Cu(H₂O)₄]²⁺.

Concept: A coordinate covalent bond forms when one species (the ligand) donates a lone pair of electrons into an empty orbital of another species (the acceptor). The acceptor must therefore have vacant, low-energy orbitals available to receive the electron pairs. Transition metal ions are good acceptors because they have vacant d (and also s and p) orbitals.

Step-by-step solution:

1. Electronic configuration of Cu: [Ar] 3d¹⁰ 4s¹
2. Cu²⁺ loses 2 electrons (one 4s and one 3d): [Ar] 3d⁹
3. The 3d⁹ configuration leaves one d-orbital with a vacancy. The ion also has empty 4s and 4p orbitals available.
4. Ligands such as NH₃ and H₂O have lone pairs. They donate these pairs into the vacant orbitals of Cu²⁺, forming coordinate bonds.

Checking the options:

• A. are coloured: Colour is a result of d-d electronic transitions. It is not the reason for bonding.
• B. are positively charged: Positive charge attracts ligands, but many positive ions (like Na⁺) do not form coordinate bonds readily. Charge alone is not the key.
• C. have vacant d-orbital: The ion can accept lone pairs into vacant orbitals. This is the standard WASSCE explanation.
• D. have unpaired electrons: Cu²⁺ does have one unpaired electron, but that explains paramagnetism, not coordinate bonding.

Answer: C

Exam tip: For "why can a metal ion form complexes or coordinate bonds?", think "vacant orbitals to accept lone pairs". For "why is it coloured or paramagnetic?", think "d-d transitions" and "unpaired electrons".