Exploring Complex Iron Salts: A Series of Analytical Chemistry Experiments

Lab # 1 The Synthesis of a Complex Iron Salt

Introduction

In this experiment, the synthesis of a complex compound containing potassium, iron, carbon, hydrogen, and oxygen is conducted. The compound's final product, characterized by emerald green crystals, is anticipated to have the empirical formula K_xFe(C₂O₄)_y.zH₂O, where zH₂O represents the water of hydration. This marks the initiation of a series of five experiments aimed at synthesizing the complex salt and determining its simplest formula.

One crucial aspect of chemical synthesis is the yield obtained compared to the theoretical amount predicted based on reaction stoichiometry.

The ratio of the actual product mass to the theoretical quantity, expressed as a percentage, is termed the 'percent yield'.

Objective

The primary objective is to prepare several grams of pure emerald green crystals of K_xFe(C₂O₄)_y.zH₂O.

Apparatus and Chemicals

Apparatus:

• Two 50-ml beakers
• One 800-ml or 1000-ml beaker
• Two 250-ml beakers
• Beaker tongs
• Electronic balance (0.001-gram accuracy)
• 9 cm watch glass
• Vacuum filtration apparatus with Buchner funnel
• Medium flow filter paper
• Black film canister with lid

Chemicals:

• FeCl₃.6H₂O solution (0.450 gram of FeCl₃.6H₂O per ml of solution)
• K₂C₂O₄.H₂O
• Ice
• Acetone

Safety, Environmental, and Economic Concerns

1. Waste solutions may be safely discarded down the drain, flushed with excess water.
2. Avoid overheating, as it may lead to violent boiling.
3. ACETONE IS A FLAMMABLE SOLVENT! Ensure all flames are extinguished and hot plates are turned off before using acetone. Cover the acetone bottle when not in use.
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Experimental Procedure

Making of the Crystals

1. Measure 10.00 ml of the stock solution of iron(III) chloride into a clean, dry 50-ml beaker.
2. Weigh 12 grams of potassium oxalate monohydrate (K₂C₂O₄.H₂O) into another clean, dry beaker.
3. Add approximately 20 ml of distilled water to dissolve the potassium oxalate monohydrate. Heat gently while stirring to dissolve the salt.
4. Pour the hot potassium oxalate solution into the beaker containing the iron(III) chloride solution and stir.
5. Cool the solution for 30-45 minutes in an ice bath to facilitate crystal formation.
6. After crystals form, carefully decant the solvent without removing any crystals.
7. Add approximately 20 ml of distilled water to the crystals and heat gently to dissolve. Decant the clear solution if residue remains undissolved.
8. Cover the beaker and set it aside to allow crystals to form undisturbed.

Day #2: Washing the Crystals

1. Clean and dry a film canister with lid. Label it with team members' names.
2. Filter the crystals using a funnel and filter paper. Wash the crystals twice with ice water and twice with acetone.
3. Spread the crystals in the bottom of a clean, dry 250-ml beaker and allow them to air dry.

Day #3: Massing and Storing the Crystals

1. Weigh the clean, dry film canister. Place the dry crystals in the canister and weigh again.
2. Store the crystals in the capped canister for future experiments. Large crystals are obtained by allowing slow, undisturbed crystal formation.

Data and Calculations

1. Volume of iron(III) chloride solution used: ____________
2. Mass of solid K₂C₂O₄.H₂O: ____________
3. Mass of film canister & complex crystals: ____________
4. Mass of film canister: ____________
5. Any other relevant data: ____________

Calculated Data

1. Mass of solid K₂C₂O₄.H₂O: ____________
2. Mass of complex crystals: ____________

 

Lab # 2 Determination of the % H2O in an Iron Oxalato Complex Salt

Introduction

In the preceding experiment, a green crystalline compound with the formula KxFe(C2O4)y.zH2O was synthesized. The forthcoming series of experiments, starting with this one, aims to determine the compound's percentage composition and derive its empirical formula. The green iron oxalato complex belongs to a category of solid chemicals known as "hydrates," which contain water chemically bound in stoichiometric amounts. Examples include Plaster of Paris (CaSO4.1/2H2O), gypsum (CaSO4.2H2O), and alum [KAl(SO4)2.12H2O]. The water of hydration in hydrates can be removed as a gas by heating the compound, leading to reactions like the following: BaCl2.2H2O(s) → BaCl2(s) + 2 H2O(g). This experiment focuses on determining the percentage of water of hydration in KxFe(C2O4)y.zH2O by heating a sample until all water of hydration is driven off.

Objective

The objective is to determine the percentage of water of hydration in the compound KxFe(C2O4)y.zH2O.

Apparatus

• Electronic top-loader balances
• Two 250-ml beakers
• Desiccator
• Beaker tongs
• Crucible tongs
• 800-ml beaker
• 150 mm watch glass
• Crucible & cover
• Drying oven

Chemicals

Student-prepared KxFe(C2O4)y.zH2O

Safety, Environmental, and Economic Concerns

1. Waste chemicals may be disposed of safely in the solid waste receptacle in the lab.
2. Precautions must be taken while handling the drying oven to prevent accidents.

Notes on Experimental Procedures

1. Measure the mass of the green crystals accurately before and after heating to calculate the mass of KxFe(C2O4)y.zH2O and the water of hydration.
2. Store crucibles & covers in a desiccator or drying oven to maintain dryness.
3. Ensure objects are at room temperature when weighing to avoid convection currents affecting balance readings.
4. Consistently use the same balance for successive weighings to ensure accuracy.

Experimental Procedure

Day 1

1. Dry the green crystals at room temperature, weigh, and store them in a capped amber bottle.
2. Place two crucibles with covers in the drying oven at 110o - 120oC for at least 25 minutes.
3. Cool the crucibles in a desiccator for 10 minutes.
4. Weigh both crucibles & covers on the analytical balance.

Day 2 and 3

1. Place approximately 1.0 gram of the green crystal into each crucible and weigh.
2. Heat the crucibles, covers, and crystal samples in the drying oven at 110o - 120oC for two hours.
3. Perform additional heating and weighing cycles until two consecutive masses agree within 0.010 gram of each other.
4. Calculate the loss in mass upon heating for each sample to determine the water of hydration percentage.
5. Discard the anhydrous crystals after the final weighings.

Data Table

Experiment Trial	Mass of Crucible, Cover, and Hydrated Crystal Sample (g)	Mass After 1st Heating (g)	Mass After 2nd Heating (g)	Mass of Empty Crucible & Cover (g)
Trial #1	__________	__________	__________	__________
Trial #2	__________	__________	__________	__________

Calculations

1. Mass of Hydrated Crystal Sample (g): __________
2. Mass of Anhydrous Crystal Sample After Final Heating (g): __________
3. Mass of Water Driven Off After Final Heating (g): __________
4. Percentage of Water of Hydration in the Crystal Sample: __________
5. Average Percentage of Water of Hydration: __________

 

Introduction

Lab #3: Standardization of a NaOH Solution

In the continuation of the green crystal lab series, the standardization of a sodium hydroxide (NaOH) solution is imperative. This experiment aims to determine the molarity of the NaOH solution accurately. Titrations, as precise and accurate analytical techniques, require meticulous attention to detail to ensure reliable results. The standardization process involves titrating a precisely weighed sample of potassium hydrogen phthalate (KHP) with NaOH, where KHP serves as the primary standard.

Primary Standard and Experimental Setup

Primary standard substances like KHP must meet specific criteria, including known purity, ease of drying, and high gram equivalent weight. KHP, with a purity of 99.95% or better, stability at 110°C, and a molecular weight of 204.22 g/mol, fits these requirements. In this experiment, KHP will neutralize NaOH in a 1:1 molar ratio.

The equipment includes solid KHP, two 250-ml Erlenmeyer flasks, a 250-ml beaker, a 50-ml buret, approximately 0.1 N NaOH solution, buret clamp, buret funnel, phenolphthalein indicator, and a ring stand.

Experimental Procedure

1. Weigh out between 0.5 - 0.8 grams of dry KHP, ensuring accurate measurements. Each group will perform four titrations, necessitating multiple weighings.
2. Add approximately 50 ml of distilled water to the Erlenmeyer flask to dissolve the KHP thoroughly.
3. While waiting for KHP to dissolve, prime the buret by rinsing it with tap water and distilled water to remove impurities.
4. Pour about 10 ml of NaOH solution into the buret using a buret funnel to remove any residual water.
5. Fill the buret slightly above the zero mark and let the base solution drain, ensuring the meniscus aligns with the 0.00-ml line.
6. Once KHP dissolves, add 5 - 7 drops of phenolphthalein indicator to the solution.
7. Titrate the KHP solution with NaOH until the first permanent pink color appears, ensuring constant stirring to facilitate the reaction.
8. Record the volume of NaOH solution used for each titration.
9. Repeat the titration with the remaining samples.

Data Table

Experiment Trial	Trial Results
	Trial #1	Trial #2	Trial #3	Trial #4
Initial mass of container (g)	___________	___________	___________	___________
Final mass of container (g)	___________	___________	___________	___________
Mass of KHP (g)	___________	___________	___________	___________
Titrations:
Final buret reading (ml)	___________	___________	___________	___________
Initial buret reading (ml)	___________	___________	___________	___________
Volume of NaOH used (ml)	___________	___________	___________	___________

Calculated Data

1. Number of moles of KHP weighed out: ___________
2. Number of moles of NaOH used to neutralize the KHP: ___________
3. Volume of NaOH used to reach the endpoint (in liters): ___________
4. Molarity of NaOH solution: ___________
5. Average molarity of NaOH solution: ___________

Conclusion

The standardization of NaOH solution using KHP as a primary standard is a crucial step in accurate volumetric analysis. The careful execution of titrations ensures reliable results for subsequent experiments. Through meticulous attention to detail and adherence to experimental protocols, the molarity of NaOH can be determined with precision and accuracy, laying the foundation for further analytical investigations.

 

Lab #4: Determination of Potassium and Iron Percentage by Ion Exchange

Introduction

In the previous experiment, the percentage of water of hydration in the green iron oxalato complex salt, KxFe(C2O4)y.zH2O, was determined. This experiment focuses on determining the percentages of potassium (K) and iron (Fe) in the salt using ion exchange chromatography. Through a titration process, the concentrations of K+ ions and Fe+ ions will be quantified in a solution passed through an ion exchange column containing a known mass of the iron complex salt.

Ion Exchange

Ion exchange resins, composed of large molecules with ionizable groups, facilitate ion exchange processes. These resins, insoluble in water and granular in nature, swell in water to form a slurry. In the presence of water, the ionizable groups undergo ionization, exemplified by the exchange of sulfonic acid groups with hydronium ions:

R-SO3-H + H2O → R-SO3-H3O+

The resin, when placed in a column and exposed to solutions containing ions, facilitates ion exchange reactions. For instance, when a potassium chloride (KCl) solution flows through a cation exchange resin column, K+ ions displace hydronium ions, resulting in the elution of aqueous HCl.

Determination of Potassium Percentage

Upon dissolution of KxFe(C2O4)y.zH2O, the salt dissociates into K+ ions and iron oxalate ions. Passing this solution through a cation exchange resin column exchanges hydronium ions for potassium ions. The eluted solution, titrated with a standardized NaOH solution, allows the determination of K+ concentration. The mass of K+ in the sample can be calculated using stoichiometry, enabling the calculation of the percentage of potassium in the salt.

Determination of Iron Percentage

After the neutralization of hydronium ions, further titration with NaOH leads to the precipitation of ferric hydroxide (Fe(OH)3). The moles of NaOH required to precipitate Fe(OH)3 and stoichiometric calculations enable the determination of the iron mass in the sample. The percentage of iron in the salt is then calculated.

The Titration Curve

Titration of the eluted solution with NaOH using a pH meter generates a titration curve with distinct endpoints: the neutralization of hydronium ions and the precipitation of ferric hydroxide. These endpoints allow for the determination of both potassium and iron percentages in the compound from a single titration curve.

Apparatus and Chemicals

The experiment requires an ion exchange column, cation exchange resin, graduated cylinder, pH paper, electronic and analytical balances, beakers, burettes, magnetic stirrer, and magnetic stir bar. Chemicals include standardized NaOH solution, buffer solutions for pH meter calibration, HCl(aq), and the green crystal sample.

Experimental Procedure

1. Mount the ion exchange column vertically and ensure it is filled with liquid above the resin.
2. Rinse the column to remove residual acid and maintain the liquid level above the resin.
3. Weigh the green crystal sample and dissolve it in water.
4. Transfer the solution to the column and collect the eluate.
5. Rinse the sample container with water and add rinses to the column.
6. Set up and prime the burette with standardized NaOH solution.
7. Titrate the eluate with NaOH, following pH changes.
8. Record the volumes of NaOH solution used for titration.

Data Table

Measurement	Value
Mass of beaker and green salt sample (g)	___________________
Initial pH of green salt solution	___________________
Volume for 1st equivalence point (V1)	___________________
Volume for 2nd equivalence point (V2)	___________________

Conclusion

The determination of potassium and iron percentages in KxFe(C2O4)y.zH2O through ion exchange chromatography and titration is an essential analytical technique. By understanding the principles of ion exchange and titration, accurate quantification of these elements is achievable. This experiment not only enhances practical laboratory skills but also provides insights into complex ion interactions and stoichiometric calculations.

 

Lab #5: Determination of the % Oxalate in an Iron Oxalato Complex Salt

Introduction

In this experiment, the objective is to determine the percentage of oxalate in the compound KxFe(C2O4)y . zH2O through titration using a standardized solution of 0.0100 M KMnO4. By titrating a solution containing a known mass of the green salt, the mass and percentage of oxalate in the sample can be accurately determined. Oxalate ions (C2O4-2) in the solution are oxidized by KMnO4 in an acidic environment to produce CO2 gas as a product. The stoichiometry of this reaction allows the calculation of the mass of C2O4-2 present in the original sample, enabling the determination of the percentage of oxalate.

Experimental Objectives

The primary objective of this experiment is to determine the percentage of oxalate ion in KxFe(C2O4)y.zH2O. This determination is crucial for understanding the composition of the green salt complex and for calculating its empirical formula and percent yield in the original synthesis reaction.

Experimental Apparatus and Chemicals

The experiment requires an electronic balance, spatula, Erlenmeyer flasks, buret funnel, buret, ring support stand, graduated cylinder, and beakers. Chemicals include student-prepared KxFe(C2O4)y.zH2O sample, 0.0100 M KMnO4 solution, concentrated H3PO4, and H2SO4.

Safety and Environmental Concerns

All waste solutions generated during the experiment can be safely disposed of down the drain and flushed with water.

Experimental Procedure

1. Weigh approximately 0.12 grams of the green salt crystal into marked 50-ml beakers.
2. Transfer the green crystals to marked 250-ml Erlenmeyer flasks and add distilled water, followed by H2SO4 and H3PO4.
3. Heat one solution to just below boiling while rinsing and filling the buret with KMnO4 solution.
4. Titrate the heated solution with KMnO4 solution and repeat the process for the second solution.
5. Record the mass of the sample, volume of KMnO4 used, and the molarity of standardized KMnO4.

Data Table

Measurement	Trial #1	Trial #2
Mass of sample (g)	_____________	_____________
Volume of KMnO4 (ml)	_____________	_____________
Molarity of KMnO4	_____________

Calculations

Calculate the moles of KMnO4 required, mass of C2O4-2 in the sample, and the percentage of C2O4-2 in the sample for each trial. Determine the average percentage of C2O4-2 in the green salt complex.

Overall Conclusions

With the percentages of key components determined, it is possible to calculate the empirical formula of the crystal. Theoretical yield and percent yield calculations should be completed, including considerations for potential discrepancies in empirical formula and percent yield. These variations may arise due to experimental errors, impurities in the sample, or incomplete reactions. These factors should be thoroughly discussed in the conclusion of the lab report.

Now that you have the percentages of the key components of your crystal, you are ready to determine the empirical formula of your crystal. Good luck! Your teacher will provide you with the accepted empirical formula and balanced equation for the reaction. Complete a theoretical yield of the product, including limiting reactant calculations, and perform a % yield calculation. Give some reasons why your empirical formula and % yield might be off.