8.2.3 Ligand Substitution Experiments

Aim:

To investigate the ligand substitution reactions of metal aqua ions with various ligands, observe the colour changes and compare the stability and rate of substitution of monodentate, bidentate, and multidentate ligands.

Introduction:

Ligand substitution involves the replacement of one or more ligands around a central metal ion by another ligand. The colour changes observed in these reactions are influenced by:

1. Type of ligand: Different ligands produce different colours when bound to the same metal ion.
2. Change in oxidation state: Oxidation states of the central metal ion can alter the electronic configuration, leading to a different colour.
3. Coordination number: Changes in the number of ligands around the metal ion also affect the observed colour.

Key Concepts:

• Monodentate ligands: Ligands that form one bond with the metal ion (e.g., $NH₃$, $H₂O$).
• Bidentate ligands: Ligands that form two bonds with the metal ion (e.g., ethylenediamine).
• Multidentate ligands: Ligands that can form multiple bonds with the metal ion (e.g., EDTA).

Materials and Equipment:

• 1.0 mol dm⁻³ metal aqua ion solutions (e.g., $Cu²⁺, Fe³⁺, Ni²⁺$)
• 1.0 mol dm⁻³ ammonia solution
• 1.0 mol dm⁻³ hydrochloric acid (HCl)
• Dropping pipettes
• Test tubes and test tube rack
• pH paper (optional)
• UV/visible spectrophotometer or colourimeter (optional for advanced analysis)
• Cuvettes (for spectrophotometer)

Method:

Part 1: Test-Tube Reactions of Metal Aqua Ions

1. Prepare the test solutions:

• Measure 2 cm³ of 1.0 mol dm$⁻³$ metal aqua ion solution (e.g., $[Cu(H₂O)₆]²⁺$) and transfer it into a test tube.

2. Addition of ammonia:

• Add 1.0 mol dm$⁻³$ ammonia solution dropwise using a dropping pipette.
• After each addition, gently shake the test tube and record any colour changes, and whether a solution or precipitate forms.
• Continue adding ammonia until the solution is in excess, and observe any further changes.

3. Addition of hydrochloric acid:

• In a separate test, add 1.0 mol dm$⁻³$ $HCl$ dropwise to another 2 cm³ of the same metal aqua ion solution.
• Shake the test tube after each addition and record the observations.

4. Repeat for other metal ions:

• Repeat the above steps for other metal aqua ions (e.g., $Fe³⁺, Ni²⁺$) and record all observations.

Observations and Expected Reactions:

Copper (II) Ions ($[Cu(H₂O)₆]²⁺$):

• With ammonia:
  • Initially, a light blue precipitate of $Cu(OH)₂$ forms.
  • Upon adding excess ammonia, the deep blue complex $[Cu(NH₃)₄(H₂O)₂]²⁺$forms.
• With hydrochloric acid:
  • The solution changes from light blue to yellow-green as $[CuCl₄]²⁻$forms.
  • The green colour is due to the presence of unreacted aqueous copper(II) ions.

Iron (III) Ions ($[Fe(H₂O)₆]³⁺$):

• With ammonia:
  • A brown precipitate of $Fe(OH)₃$ forms, indicating hydrolysis of the aqua ion.
• With hydrochloric acid:
  • A yellow solution forms due to the formation of $[FeCl₄]⁻$

Part 2: Comparing the Stability of Ligand Complexes

Bidentate and Multidentate Ligands:

• Complexes formed with bidentate and multidentate ligands are more stable than those formed with monodentate ligands.
• This is due to the chelate effect, where each multidentate ligand displaces multiple water molecules, leading to an increase in entropy and making the formation of these complexes more favourable.

Entropy and the Chelate Effect:

• When a multidentate ligand binds to a metal ion, the number of species in the system increases, raising entropy.
• The higher entropy makes the formation of the complex more thermodynamically favourable, resulting in greater stability of complexes involving multidentate ligands.

Part 3: Measuring the Rate of Ligand Substitution (Optional Advanced Experiment)

Method

1. Set up the UV/Visible Spectrophotometer:

• Measure the initial concentration of the metal aqua ion solution using a UV/visible spectrophotometer or a simpler colorimeter.

2. Monitor the ligand substitution:

• Add a monodentate ligand (e.g., $Cl⁻$or $NH₃$) to the metal aqua ion solution.
• Every 30 seconds, extract 0.1 cm³ of the reaction mixture, dilute it with water, and place it in a cuvette for analysis.
• Measure the absorbance at the wavelength of maximum absorption for the complex.

3. Plot absorbance against time:

• Plot a graph of time (x-axis) vs absorbance (y-axis).
• Repeat the same process with a multidentate ligand (e.g., EDTA) and compare the rates of substitution for monodentate and multidentate ligands.

Calculation:

The absorbance data follows the Lambert-Beer law:

$$A = ɛcl$$

Where:

• A = absorbance
• ɛ = molar extinction coefficient
• c = concentration
• l = path length of the cuvette (typically 1 cm) Using the absorbance data, calculate the rate of substitution for both monodentate and multidentate ligands.

Conclusion:

• Ligand substitution reactions vary depending on the type of ligand, oxidation state, and coordination number.
• Monodentate ligands (such as $NH₃$ or $Cl⁻$) may show fast substitution, but complexes involving bidentate or multidentate ligands are more stable due to the chelate effect.
• The rate of substitution can be quantitatively measured using UV/visible spectrophotometry, comparing how fast monodentate ligands are replaced versus multidentate ligands.

Table showing factors affecting colour:

Factor	Complex Ion	Colour	Shape/Coordination
Changing the ligand, but not the coordination number	$[Cu(H₂O)₆]²⁺$	Blue	Octahedral
	$[Cu(NH₃)₄(H₂O)₂]²⁺$	Deep blue	Octahedral
Changing the ligand and the coordination number	$[Cu(H₂O)₄]²⁺$	Blue	Octahedral
	$[CuCl₄]²⁻$	Yellow	Tetrahedral
Changing the oxidation state	$[Fe(H₂O)₆]²⁺$	Green	Octahedral
	$[Fe(H₂O)₆]³⁺$	Pale violet	Octahedral