Complexometric Titration: Coordination Complexes

• Complexometric titrations are discussed in the context of lectures 8 and 9.

Co-ordination Complexes

• Co-ordination complexes are neutral or ionic compounds.
• Coordinate bond formation links a metal ion (cation, electron acceptor) with a complexing agent (electron donor).
• Ionic bonds involve complete electron transfer, forming ions, followed by electrostatic attraction, seen in NaCl.
• Covalent bonds involve shared electrons between atoms, exemplified by Cl₂.
• The reaction creating a complex can be seen as a Lewis acid-base reaction.
• The central metal atom acting as a Lewis acid (electron acceptor) and the ligand as a Lewis base (electron donor).

Complexation Reaction

• Complexation reactions form a complex from a central metal atom (M) and one or more ligand (L) molecules; n = coordination number.
• Coordination number indicates the maximum bonds a central metal ion can form.
• Coordination numbers commonly range from 2 to 8, with 4 and 6 being most common.
• The coordination number often equals double the metal's valency, like Cu²⁺ having a coordination number of 4.
• Some ions can display multiple coordination numbers.

Ligands

• Ligands are classified based on coordinating atoms.
• Monodentate ligands bind at one point, donating a lone electron pair creating one coordinate bond.
• Monodentate ligands include anions (N, O, S, F⁻, Cl⁻, Br⁻, I⁻, CN⁻, SCN⁻) and molecules (H₂O & NH₃), such as Ag(NH₃)₂⁺ and Cu(NH₃)₄²⁺.
• Bidentate ligands contain two coordinating atoms, such as ethylene diamine.
• Tris(ethylene-diamine) cobalt(III) ([Co(en)₃]³⁺) forms a 6 coordinate complex.
• Multidentate ligands contain more than two coordinating atoms, for instance, ethylenediamine tetraacetic acid with 2 nitrogen and 4 oxygen atoms (hexadentate).

Stability of Complexes

• The stability reflects how readily a complex is formed under equilibrium conditions.
• Stability increases with a higher log Kst value.
• Transition metals strongly attract ligands, forming stable complexes.
• Stability is enhanced by smaller ionic radii of the metal.
• Zn²⁺ (i.r.= 69 picometer, log K = 18.6) forms more stable complexes than Mn²⁺ (i.r.= 71 picometer, log K = 13.8).
• High electric charge on the ion increases complex stability.
• Fe(III)-EDTA (log K = 25.3) is more stable than Fe(II)-EDTA (log K = 14.6).
• Ligand complexation tendency decreases in the order C>N>O>S>F.
• Larger anion radii favor stability.
• Anion stability order: PbI₄²⁻ > PbBr₄²⁻ > PbCl₄²⁻ > PbF₄²⁻.
• Higher temperatures reduce complex stability by increasing ionization and decreasing Log K.
• Ionization suppressors, like ethanol, enhance stability.
• Chelate effect contributes to stability.

Chelate Effect

• Multidentate ligands create more stable metal complexes than similar monodentate ligands.
• This effect results from forming a 5-membered "ring" structure with the metal and two ligand atoms.
• Complex formation examples include Cd²⁺ with methylamine and ethylenediamine, showing stability constants.

Complexones

• Complexones, such as aminopolycarboxylic acids, are complexing agents.
• Complexone II (EDTA) and Complexone III (Na₂EDTA) both have 6 donor groups.

EDTA Complexes

• Na₂EDTA (Na₂H₂Y): In aqueous medium, dissociates into a complex-forming ion (H₂Y²⁻) which reacts with metals in a 1:1 ratio.
• EDTA complexes with metal ions having a charge number 2 are stable in alkaline or slightly acidic solutions, while complexes with ions having a charge number 3 or 4 can exist in solutions with higher acidity.
• pH requirements vary for complex stability with selected metals.
• At pH 1-3: Zr⁴⁺, Hf⁴⁺, Th⁴⁺, Bi³⁺, Fe³⁺
• At pH 4-6: Pb²⁺, Cu²⁺, Zn²⁺, Co²⁺, Ni²⁺, Mn²⁺, Fe²⁺, Al³⁺, Cd²⁺
• At pH 8-10: Ca²⁺, Sr²⁺, Ba²⁺, Mg²⁺.

Stability Constants

• These constants, or formation constants (K), measure the stability of the complex, shown as: Mn+ + Y4− ⇄ MY(n–4)+, K = [MY(n–4)+] / [Mn+ ] [Y4−].

EDTA Titration Curve

• In EDTA titrations, three regions exist: excess metal, equivalence point, and over-titration/excess EDTA.
• Region 1 illustrates excess metal.
• Region 2 occurs at the equivalence point.
• Region 3 involves over-titration or excess EDTA.
• At 0% titration, the pNi equals 2.00.
• Between 0% and 100% titration, the amount of free Ni(II) can be calculated with an equation.
• At 100% titration, KMY equals 1.87 x 10¹⁸.

Metallochromic Indicators

• They are compounds that change color upon binding to a metal ion.
• Mg-In + EDTA → Mg-EDTA + In (red to colourless to blue).
• Stability of M-In complex must be less than the stability of M -EDTA complex.
• A metal-indicator complex must possess sufficient stability and produce a sharp color change.
• Indicators must be sensitive to metal ions, changing color near the equivalence point.
• The indicator must fulfill the requirements within the titration's pH range.
• Before titration: M + In = free M + In-M (color 1).
• Free M + EDTA = EDTA-M.
• Excess EDTA + In-M = EDTA-M + In- (color 2).
• EDTA reacts more readily with uncomplexed metal.
• M-Ind is harder for EDTA to react with so we must insure that only a small amount of indicator is used.
• Murexide It is the first metallochromic indicator used in EDTA titrations.
• Murexide is mainly used for calcium titration at pH 12 (In--).
• Eriochrome Black It is the most widely used metal indicator for EDTA titrations.
• Eriochrome Black is suitable for most metals eg. Mg 2+, Zn 2+, Cd 2+ at alkaline pH.