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IB DP Chemistry -R3.4.8 Ligands and coordination bond – Study Notes

IB DP Chemistry -R3.4.8 Ligands and coordination bond - Study Notes - New Syllabus - 2026, 2027 & 2028

IB DP Chemistry – R3.4.8 Ligands and coordination bond – Study Notes – New Syllabus

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Reactivity 3.4.8 - Coordination Bonds and Ligands in Complex Ions

Reactivity 3.4.8 – Coordination Bonds and Ligands in Complex Ions

Coordination bonds are a special type of covalent bond formed when a ligand donates a lone pair of electrons to a transition metal cation, resulting in the formation of a complex ion.

Definitions

  • Ligand – A species that has a lone pair of electrons and can form a coordinate bond with a metal ion. Ligands act as Lewis bases.
  • Transition metal cation – A positively charged transition metal ion that can accept electron pairs. These ions act as Lewis acids.
  • Coordinate bond – A covalent bond in which both bonding electrons are donated by the same atom (the ligand).
  • Complex ion – A central metal ion bonded to one or more ligands via coordinate bonds, forming a charged species.

Nature of Transition Metal Cations

  • Transition metals have high charge densities and unfilled d-orbitals that can accept electron pairs from ligands.
  • This allows them to form stable complexes with a wide variety of ligands.

Nature of Ligands

Ligands are Lewis bases – they donate a lone pair to form a coordinate bond.

Common ligands include:

    • Monodentate (donates one pair): \( \text{NH}_3 \), \( \text{H}_2\text{O} \), \( \text{Cl}^- \), \( \text{CN}^- \)
    • Bidentate (donates two pairs): Ethane-1,2-diamine (en), oxalate \( (\text{C}_2\text{O}_4^{2-}) \)
    • Polydentate (multiple pairs): EDTA\( ^{4-} \)

General Formula of a Complex Ion

\( [\text{Metal}^{n+}(\text{ligand})_x]^{\text{charge}} \)

  • The square brackets enclose the entire coordination complex.
  • The charge on the complex depends on the charge of the metal ion and the total charges on all ligands.

Example

Formation of the tetraamminecopper(II) complex:

\( \text{Cu}^{2+} + 4\text{NH}_3 \rightarrow [\text{Cu(NH}_3)_4]^{2+} \)

▶️Answer/Explanation
  • \( \text{Cu}^{2+} \) is the central metal ion (Lewis acid).
  • Each \( \text{NH}_3 \) ligand donates a lone pair to form a coordinate bond with Cu\( ^{2+} \).
  • 4 coordination bonds are formed in total, resulting in a stable octahedral or square planar structure depending on the metal and ligands involved.

Example

Formation of the hexaaquairon(III) complex:

\( \text{Fe}^{3+} + 6\text{H}_2\text{O} \rightarrow [\text{Fe(H}_2\text{O})_6]^{3+} \)

▶️Answer/Explanation
  • \( \text{Fe}^{3+} \) is the central transition metal ion.
  • Each \( \text{H}_2\text{O} \) molecule acts as a monodentate ligand, donating a lone pair on oxygen.
  • 6 coordinate bonds are formed, resulting in an octahedral complex.

Importance of Coordination Complexes

  • Coordination compounds are crucial in biological systems (e.g. hemoglobin contains iron-porphyrin complexes).
  • They are widely used in catalysis, medicine (e.g. cisplatin), and analytical chemistry.

Deducing the Charge on a Complex Ion

The charge on a complex ion can be determined by adding the oxidation state (charge) of the central metal ion and the total charges of all the ligands attached to it.

General Formula

Let:

  • M be the central metal ion with charge \( x \)
  • L be the ligand with charge \( y \), and there are \( n \) ligands

Then:

Total charge of complex ion = \( x + n \times y \)

Step-by-step Method

  1. Determine the oxidation state of the central metal ion (usually provided or deduced from context).
  2. Multiply the charge of each ligand by the number of such ligands present.
  3. Add the metal ion charge and the total ligand charges.
  4. The result is the charge on the overall complex ion.

Common Ligands and Their Charges

  • \( \text{H}_2\text{O} \), \( \text{NH}_3 \) – neutral ligands, charge = 0
  • \( \text{Cl}^- \), \( \text{CN}^- \), \( \text{OH}^- \) – monodentate anionic ligands, charge = -1
  • \( \text{C}_2\text{O}_4^{2-} \) (oxalate), \( \text{EDTA}^{4-} \) – multidentate anionic ligands

Example

Deduce the charge on the complex ion \( [\text{Cr(NH}_3)_6] \)

▶️Answer/Explanation
  • Chromium is in the +3 oxidation state → \( \text{Cr}^{3+} \)
  • \( \text{NH}_3 \) is a neutral ligand → charge = 0
  • Total ligand charge = \( 6 \times 0 = 0 \)
  • Charge on complex = \( +3 + 0 = +3 \)

Final answer: \( [\text{Cr(NH}_3)_6]^{3+} \)

Example

Deduce the charge on the complex ion \( [\text{Fe(CN)}_6] \)

▶️Answer/Explanation
  • Assume iron is in the +3 oxidation state → \( \text{Fe}^{3+} \)
  • Each \( \text{CN}^- \) ligand has a charge of -1
  • Total ligand charge = \( 6 \times -1 = -6 \)
  • Charge on complex = \( +3 + (-6) = -3 \)

Final answer: \( [\text{Fe(CN)}_6]^{3-} \)

Example

Deduce the charge on the complex ion \( [\text{CuCl}_4] \)

▶️Answer/Explanation
  • Assume copper is in the +2 oxidation state → \( \text{Cu}^{2+} \)
  • Each \( \text{Cl}^- \) ligand has a charge of -1
  • Total ligand charge = \( 4 \times -1 = -4 \)
  • Charge on complex = \( +2 + (-4) = -2 \)

Final answer: \( [\text{CuCl}_4]^{2-} \)

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