### Introduction In this simple experiment we will study the complex formation of Fe3+ and salicylic — Physical Chemistry — Kinetics Chemistry Question

### Introduction
In this simple experiment we will study the complex formation of Fe³⁺ and salicylic acid in the aqueous solution. The empirical formula of the complex will be determined and also its stability constant can be estimated.
Several stable complexes between ferric ion and salicylic acid H₂Sal have been known. Their structures and compositions are much dependent on pH. In acidic solution, a violet complex is formed. At neutral pH, a different dark-red complex forms, and in basic solution the complex that forms is orange. This experiment will be carried out at pH of about 2. Under this condition, the hydrolysis of ferric ion is largely suppressed. To simplify the calculations, we will not consider to the dissociation of H₂Sal during the complex formation. Thus, regardless of the structure of the complex, we can present the complex formation equilibrium as:
Fe³⁺ + n H₂Sal <=> Fe³⁺(H₂Sal)n

Thus, the stability constant Kf is defined as:
Kf = [Fe³⁺(H₂Sal)n] / ([Fe³⁺] [H₂Sal]^n) (1)

where the [Fe³⁺] and [H₂Sal] refer to the concentrations of the free species.
The complex Fe³⁺(H₂Sal)n absorbs most strongly at 528 nm (neither Fe³⁺ nor H₂Sal absorb at this wavelength). Its concentration is related to the optical absorbance through Beer’s law, which is:
A = ε l [Fe³⁺(H₂Sal)n]
where ε is the molar extinction coefficient for the complex and l is the optical path length.
Job’s method can be used to find the empirical formula of the complex. Following this method, equimolar solutions of Fe³⁺ and H₂Sal are prepared, and then mixed in ratios of 1:9; 2:8 … 9:1. The total reagent concentrations therefore are the same in each solution. Maximum amount of equilibrium complex will be formed when the proportions of reagents employed correspond to the empirical formula of the complex and can be deduced through the measurement of optical absorbance.

### Chemicals and Reagents
* A solution of 0.0025 mol dm^-3 Fe³⁺ made by dissolving the appropriate amount of ferric ammonium sulfate in 500 cm₃ of 0.0025 mol dm^-3 sulfuric acid.
* A solution of 0.0025 mol dm^-3 salicylic acid made by dissolving the appropriate amount of salicylic acid in 500 cm₃ 0.0025 mol dm^-3 sulfuric acid.
* Saturated solution of salicylic acid (about 50 cm₃) in 0.0025 mol dm^-3 sulfuric acid.

### Apparatuses and Glassware
* Glass beaker: 100 cm₃, 50 cm₃
* Burette: 25 cm₃
* Volumetric flask: 500 cm₃
* Wash bottle
* Electronic balance with readability of 0.0001 g
* UV-vis spectrophotometer
* Glass cuvettes.

### Experimental procedure
Step 1. Determine the empirical formula of the complex by Job’s-method
1. Prepare in 100 cm₃ beakers (should be dry and clean) a series of nine mixtures of the 0.0025 mol dm^-3 iron(III) and the 0.0025 mol dm^-3 salicylic acid solutions, plus 10.0 cm₃ 0.0025 mol dm^-3:
[VISUAL]
(Note: Use burette to measure the volumes of the solution.)
2. Measure the absorbance of each mixture.
3. Plot absorbance versus volume of Fe³⁺. Absorbance should be highest for the stoichiometric mixture.

Step 2. Determine the molar extinction coefficient ε of the complex
1. Pipette out 1.00, 2.00, 3.00, 4.00, 5.00, 6.00 cm₃ of 0.0025 mol dm^-3 iron(III) solution into 5 beakers (100 cm₃). To each beaker add 10.00 cm₃ of saturated salicylic acid solution and enough HCl solution (0.0025 mol dm^-3) to reach the total volume to 20.00 cm₃.
2. Measure the absorbance of each solution.
3. Plot absorbance versus [Fe³⁺] (Because the salicylic acid is in excess, it is assumed that the concentration of iron equals the concentration of the complex).
4. Calculate ε from the linear plot.

Step 3. Determine the stability constant Kf
1. Prepare (in 100 cm₃ beakers) three mixtures of the same volumes of 0.0025 mol dm^-3 Iron(III) and the 0.0025 mol dm^-3 salicylic acid solutions and plus 0.0025 mol dm^-3 HCl solution to total volume of 20 cm₃:
[VISUAL]
2. Measure the absorbance of each solution.
3. Calculate the initial concentration of Fe(III) and H₂Sal in each solution.
4. From the measured absorbance and observed ε value determined in step 2(4)). Calculate the concentration of the complex in each solution.
5. Calculate the equilibrium concentration of Fe(III) and H₂Sal. Assume that:
[Fe³⁺]eq = [Fe³⁺]initial – [Fe³⁺(H₂Sal)n]
[H₂Sal]eq = [H₂Sal]initial – n × [Fe³⁺(H₂Sal)n]
6. Calculate the equilibrium constants Keq for each solution (using equation 1) and determine an average value.