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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 KxFe(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'.
The primary objective is to prepare several grams of pure emerald green crystals of KxFe(C2O4)y.zH2O.
Apparatus:
Chemicals:
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Cover the acetone bottle when not in use.
Heat gently while stirring to dissolve the salt.
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.
The objective is to determine the percentage of water of hydration in the compound KxFe(C2O4)y.zH2O.
Student-prepared KxFe(C2O4)y.zH2O
| 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 | __________ | __________ | __________ | __________ |
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 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.
| 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) | ___________ | ___________ | ___________ | ___________ |
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.
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 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.
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.
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.
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.
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.
| 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) | ___________________ |
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.
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.
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.
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.
All waste solutions generated during the experiment can be safely disposed of down the drain and flushed with water.
| Measurement | Trial #1 | Trial #2 |
|---|---|---|
| Mass of sample (g) | _____________ | _____________ |
| Volume of KMnO4 (ml) | _____________ | _____________ |
| Molarity of KMnO4 | _____________ | |
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.
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.
Exploring Complex Iron Salts: A Series of Analytical Chemistry Experiments. (2024, Feb 23). Retrieved from https://studymoose.com/document/exploring-complex-iron-salts-a-series-of-analytical-chemistry-experiments
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