Preparation of Iron Sulfate

The objective of this lab is to prepare iron sulfate by allowing iron metal to react with sulfuric acid. Through this experiment, students will gain practical experience in using tools and equipment necessary for scientific analysis and research, demonstrating safe practices in the chemical laboratory, understanding electrons in atoms and the periodic table, chemical bonding, chemical reaction and quantities in chemical reactions, molecules, compounds, and chemical composition. It also emphasizes the proper use of exponential notation and significant figures, recording investigation results, using scientific reasoning to evaluate phenomena, and identifying the unifying themes in the scientific field of study.

In the discussion, it is explained that some compounds are not naturally occurring and must be synthesized through chemical reactions. One method involves a metal reacting with an aqueous solution of a strong acid. In this specific experiment, iron metal reacts with sulfuric acid to form iron sulfate.

The procedure involves weighing a 125 mL Erlenmeyer flask and adding iron powder to it. The flask is then reweighed, and approximately 25 mL of 3 M sulfuric acid is carefully added to the flask. The reaction flask is placed in a water bath, which is heated to boiling using a Bunsen burner. After the reaction flask is soaked in the hot water bath, the contents are gravity filtered to remove any unreacted iron. The filtrate is collected, cooled in an ice bath, and crystals should form. If crystallization does not occur, a seed crystal is provided. The crystals are then filtered using a suction filtration apparatus, and the solid precipitate is washed with ethyl alcohol. The crystals are spread out on a labeled filter paper and allowed to dry. The weight of the dried crystals is determined, and the theoretical yield and percentage yield are calculated assuming the product can be either Fe2(S04)3 ∙ 9H20 or FeSO4 ∙ 7H20.

In addition to the iron sulfate preparation, students are instructed to conduct the HYDRATES Lab, Parts A and B, using the dry iron sulfate prepared in this lab. The formula of the hydrated iron sulfate is determined according to the directions in the HYDRATES Lab, Part B. Two formulas for the iron sulfate hydrate are calculated assuming the presence of either Fe2(S04)3 ∙ 9H20 or FeSO4 ∙ 7H20. Based on the data, students decide which oxidation state of iron is most likely to be present.

Through this lab, students gain practical experience in conducting chemical reactions, performing filtration, calculating yields, and analyzing the composition of compounds. The lab also highlights the importance of accurate measurements, understanding reaction stoichiometry, and the concept of hydrates in chemistry.

Objective

·         Use the tools and equipment necessary for basic scientific analysis and research

·         Demonstrate safe practices in the Chemical Laboratory

·         Demonstrate an understanding of Electrons in Atoms and the Periodic Table and Chemical Bonding

·         Demonstrate an understanding of Chemical Reaction and Quantities in Chemical Reactions

·         Demonstrate an understanding of Molecules, Compounds and Chemical Composition

·         Demonstrate the proper use of Exponential Notation and Significant Figures

·         Record the results of investigation through writing

·         Use scientific reasoning to evaluate physical and natural phenomena

·         Identify the unifying themes of the scientific field of study

Materials

·         125 mL Erlenmeyer Flask

·         400 mL Beaker

·         Iron Powder

·         3M Sulfuric acid (H2S04)

·         Boiling Chips

·         Bunsen Burner

·         Reaction Flask

·         Liquid Funnel

·         Filter Paper

·         Ice Bath

·         Suction Filter

Discussion

Certain compounds exist naturally as minerals or are commonly found in rocks. Examples include SiO2, which is present in sand, NaCl, found in halite, and CaCO3, which constitutes limestone. However, many other compounds do not occur naturally and must be synthesized from readily available substances through chemical reactions. One method of synthesis involves the reaction of a metal with an aqueous solution of a strong acid. For instance,

Zn + 2 H30+  →  Zn2+ + 2 H20 + H2

It is important to note that the reaction also produces hydrogen gas. If hydrochloric acid is used, it results in a solution of zinc chloride:

Zn + 2 H30+ + 2 Cl- →  Zn2+ + 2 Cl- + 2 H20 + H2

By evaporating the water, solid zinc chloride, ZnCl2, can be obtained. The salt may retain one or more water molecules as water of hydration, as described in the "HYDRATES" Lab. In this particular experiment, iron metal will react with sulfuric acid to produce iron sulfate.

Procedure

1.       Perform the following steps in the HYDRATES Lab:

2.       Measure the weight of a 125 mL Erlenmeyer flask using a top-loading balance, rounding to the nearest 0.01 g.

3.       Add approximately 3.00 g of iron powder to the flask and reweigh it.

4.       Label the flask with your initials and place it in the hood.

5.       Take caution and add 25 mL of 3 M H2S04 to the flask. (CAUTION: Sulfuric acid is a strong acid. In case of spills, immediately flush the affected area with cold water and inform the instructor.)

6.       Once the reaction flask is safely inside the hood, prepare a water bath. Pour 200 mL of tap water into a 400 ml beaker and add 4 or 5 boiling stones. Heat the water with a Bunsen burner until it reaches boiling point. Once boiling, remove the burner and extinguish the flame.

7.       Carefully place the hot water bath inside the hood and gently insert your reaction flask into the bath. Let it remain submerged in the hot water for 20-30 minutes.

8.       After removing the flask from the water bath, use a gravity filter to separate the contents and eliminate any unreacted iron. Collect the filtrate in a clean Erlenmeyer flask.

9.       Cool the flask and its contents in an ice bath for 15 minutes to encourage crystallization. If crystals do not form, seek assistance from the instructor for a seed crystal.

10.    Utilize the provided suction filtration apparatus to filter the crystals and wash the solid precipitate with 5 mL of ethyl alcohol.

11.    Following filtration, spread the crystals on a labeled filter paper that has been weighed to the nearest 0.01 g using the top-loading balance. Allow the crystals to dry in this manner.

12.    Once the crystals have dried completely, measure the weight of the product.

13.    Calculate the theoretical yield and percentage yield, considering two potential compounds:

Fe2 (S04)3 ∙ 9H20 or FeSO4 ∙ 7H20

HYDRATES Lab:

1.       Conduct Parts A and B of the HYDRATES Lab. Instead of using an unknown substance, employ approximately 2 grams of the dry iron sulfate prepared in the previous steps.

2.       Follow the directions outlined in HYDRATES Lab, Part B to determine the formula of the hydrated iron sulfate.

3.       Complete the Report Sheet by recording your observations and calculations.

4.       Calculate two possible formulas for the iron sulfate hydrate: Fe2(S04)3 ∙ 9H20 or FeSO4 ∙ 7H20.

5.       Analyze the data to determine which oxidation state of iron is most likely to be present.

Report Sheet for PREPARATION of IRON SULFATE

Name_______________________________________                          Date___________________________

Lab / Section_________________________________

Theoretical Yield Calculations:

Iron(III) Sulfate Nonahydrate:

(F.W.=562; anhydrous F.W. = 400)

Iron(II) sulfate Heptahydrate:

(F.W. = 278; anhydrous F.W.=152)

Percent Actual yield Calculations:

Iron(III) sulfate Nonahydrate:

Iron(II) Sulfate Heptahydrate:

Review Questions

1.       What is the objective of the experiment?

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2.       Name two tools or equipment used in the lab.

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3.       Describe the process of gravity filtration and its purpose in the experiment.

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4.       How does the reaction between iron metal and sulfuric acid produce iron sulfate?

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5.       Explain the concept of theoretical yield and how it is calculated in this experiment.

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6.       What is the purpose of using an ice bath during the crystallization step?

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7.       What is the difference between Fe2(S04)3 ∙ 9H20 and FeSO4 ∙ 7H20?

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8.       How is the weight of the dried iron sulfate crystals determined?

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9.       What are some safety precautions that should be followed when working with sulfuric acid?

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10.    Describe the role of the water bath in the experiment and why it is heated.

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11.    What is the significance of using a Bunsen burner in this lab?

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12.    Explain the concept of a hydrate and its relevance to the experiment.

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13.    How does the experimental process in this lab relate to the law of conservation of mass?

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14.    Discuss the importance of accurate measurements and weighing in this experiment.

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15.    How does the percentage yield provide information about the efficiency of the reaction?

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16.    Describe the purpose of using boiling chips in the water bath.

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17.    What are some potential sources of error in this experiment and how can they be minimized?

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18.    Explain how the concept of atomic weight is related to determining the equivalent weight of a metal.

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19.    Discuss the role of the filtration apparatus and why it is necessary in the experiment.

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20.    How does the preparation of iron sulfate in this lab contribute to our understanding of chemical composition and synthesis?

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