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13 Dec 2019

Question 1: Hydrogen peroxide will be used to oxidize the iron(II) in the complex ion to iron(III). Several things happen in this reaction:

Iron(II) is oxidized to iron(III)

The oxalate ions that were attached to the iron(II) in the complex ion return to the solution (as oxalate ions)

Each hydrogen peroxide molecule is reduced to form two hydroxide ions

Part A - Oxidation of iron(II) to iron(III).

Write a balanced net ionic equation for the oxidation of the iron(II) complex ion with hydrogen peroxide. Make certain that your equation depicts all three things described in the introduction and that is it blanced for nujmber of electrons transferred (charge) as well as for atoms.

Question 2:Part A - The dissolution of iron(II) ammonium sulfate

In the first step of experiment #9, iron(II) ammonium sulafte is desolved in water. Write the equation for this process.

Write the equation for the dissolution of Fe(NH4)2(SO4)2. Include the phases in your equation

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Lelia LubowitzLv2

17 Dec 2019

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you will synthesize iron (II) oxalate, a yellow precipitate, byreacting iron (II) ammonium sulfate hexagydrate,Fe(NH4)2(SO4)2.6H2O, and saturated oxalic acid solution in abeaker. The compound containing the iron (II) ion will be dissovedin water.

1)List the predominant species in the Fe(NH4)2(SO4)2.6H2Osolution

2)write the net ionic equation for the formation of iron (II)oxalate via the reaction of potassium oxalate andFe(NH4)2(SO4)2.6H2O

3)Potassium oxalate solutionwill be added to the solid toprovide the ligand for the complex ion. Is the cation or the anionthe ligand? The formula of the ligand?

4)The oxidation number of the iron in the complex ion is +3. Isan oxidizer or a reducer required to change the oxidationnumber ofthe iron?

5)Write the balanced net ionic equation for the reaction ofaqueous iron (II) ion reacting with H2O2 in acidic medium

coffeellama276

INTRODUCTION

Many transition metal cations can have more than one possible charge (Worksheet 1). Differently charged species of the same element can have different chemical and physical properties. Knowing the charge can allow separation of the different ions based on their properties. For example, Cu²⁺ is blue and stable in solution, while Cu⁺ is colorless and reacts with oxygen. Hemoglobin contains iron as Fe²⁺, but it is inactive at binding oxygen when the iron gets oxidized to Fe³⁺. There are many similar situations in which it is important to always properly indicate charges of ions.

This laboratory experiment uses the difference in chemical properties of Fe²⁺, iron(II), and Fe³⁺, iron(III), to quantify the amount of each present in a solution of unknown concentrations of each. You will first determine the amount of iron(II) in your sample by titrating with a reagent that does not react with iron(III). You will then titrate a second sample after converting all the iron(III) to iron(II) in order to determine the total iron present. The percent composition of Fe²⁺ in the unknown mixture will be calculated.

BACKGROUND

The concentrations of both iron(II) and iron(III) ions in a solution can be determined by titration with potassium permanganate. The reaction is an oxidation-reduction reaction, or redox, for short. A brief introduction to oxidation-reduction reactions is given in the background of the âAcids and Basesâ experiment. The permanganate ion effectively oxidizes iron(II) to iron(III) in an acidic solution while being reduced to manganese(II) ion. The balanced net ionic reaction is:

5 Fe²⁺(aq) + MnO4^â(aq) + 8 H⁺(aq) â 5 Fe³⁺(aq) + Mn²⁺(aq) + 4 H₂O(l)

Iron(III) is not oxidized by the permanganate ion, and thus does not react. Thus in order to determine the concentration of iron(III) ions, they must first be reduced to iron(II) ions. Zinc metal is used as the reducing agent, also under acidic conditions.

Zn(s) + 2 Fe³⁺(aq) â 2 Fe²⁺(aq) + Zn²⁺(aq) (reduction of the Fe3+)

Zn(s) + 2 H⁺(aq) â Zn²⁺(aq) + H₂(g) (a side reaction that oxidizes excess Zn)

The Zn2+ formed does not interfere with the subsequent titration because Zn2+ is not oxidized by permanganate. Notice that converting all the iron(III) to iron(II) means that the total iron concentration ([Fe2+] + [Fe3+]) is titrated, not just the iron(II) originally present.

Techniques

This is a lab will utilize several of the techniques introduced in the experiment âLaboratory Techniques and Measurements,â the procedures for which are given in Appendix B. Poor laboratory technique and misuse of the glassware will result in poor data results, which will affect the grade you receive on the laboratory report. In order to use your time efficiently in lab, review how to properly:

rinse glassware prepare a standard solution use a volumetric pipette

weigh by difference set up and use a burette use a graduated cylinder

heat samples on a hot plate

Analysis of Iron by Oxidation-Reduction Titration

This lab uses the common quantitative technique of titration. The only difference between the titration you will be doing in this lab and that done in the Molar Mass of a Known Acid lab is the type of reaction used: this lab will use a redox reaction in place of the acid-base reaction performed in the previous experiment.

In order to determine the concentrations of iron(II) and iron(III) ions in a solution containing both ions, two portions of the same sample must be titrated separately. This first titration will determine only the iron(II) concentration. A second sample will be treated with zinc metal in order to reduce the iron(III) to iron(II). The resulting solution will be titrated with potassium permanganate. The result of this titration will give the total iron concentration in the sample. The difference between the first and second titration is the equal to the iron(III) concentration.

In this titration, potassium permanganate serves as its own indicator. The intense purple color of the permanganate solution becomes colorless when the permanganate is reduced to manganese(II). Thus, as iron(II) is oxidized with the addition of the permanganate solution from the burette it will impart a purple color that will fade. However, once there is no iron(II) left in solution a faint purple color will remain, marking the endpoint of the titration. This color can be slightly altered by the presence of iron(III), which is yellow in solution. Remember that iron(III) is generated by the reaction between iron(II) with permanganate, increasing the yellow appearance of the solution and which can interfere with the detection of the endpoint. Phosphoric acid is added to help minimize this effect.

pre lab questions: