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12 Dec 2019

I've only gotten five correct so far:

* iodine and bromide ion = no reaction
* iron(III) with bromide ion = no reaction
* hydrogen in copper(II) ion = no reaction
* bromine and iodide ion = reaction visible
* iron(III) with iodide ion = reaction visible

What about in these reactions?:

* copper and zinc ion = ?
* zinc and copper(II) ion = ?
* iron and copper(II) ion = ?
* iron and zinc ion = ?
* copper and iron(II) ion = ?
* zinc and iron(II) ion = ?

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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:

Nomenclature lab : will rate answer

Nomenclature lab: will choose best answr

= nonmetal

-3

-2

-1

= metalloid

Type I Binary Ionic Compounds

Type I binary ionic compounds contain a metal and a nonmetal AND the metal that is present only forms one type of cation. Metals with only one cation(shaded below with charges). Both the metal and the nonmetal form ions, which is why it is called an ionic compound.

From the following list, cross out those compounds that do NOT belong in the category for Type I binary ionic compounds.

NaCl FeCl2 CaCl2 TiO2 MgO AlBr3 KCl K2S BeF2 Cu2O3 AgCl Zn3N2

Formula and name examples for Type I binary ionic compounds:

KI = potassium iodide BaO = barium oxide ZnF2 = zinc fluoride Na2S = sodium sulfide

Ag3N = silver nitride BeCl2 = beryllium chloride

What type of element is always listed first (metal or nonmetal)? ____________ second? ___________

Is the name of the first element in the compound different from the element? (yes/no)

What is the common ending for all the names? ______________

In zinc fluoride, there are 2 fluoride atoms, are they indicated in the name? (yes/no)

What is the charge on the zinc ion? _______

What is the charge on the fluoride ion? _______

Why do you need one zinc ion and two fluoride ions for the formula for zinc fluoride?

Why do you need two sodium ions for every sulfide ion in sodium sulfide?

Try not to look in your book and determine the rules for naming type I binary ionic compound when given the formula.

Try not to look in your book and determine the rules for writing the formula for a type I binary compound when given the name.

Name each of the type I binary ionic compounds listed in Question 1 of the Type I section in words.

Type II Binary Ionic Compounds

Type II binary ionic compounds also contain a metal and a nonmetal however the metal that is present here can form more than one type of cation. Metals with multiple possible charges are listed in the periodic table as blank. Type II metals are NOT Type I metals. Again, both the metal and the nonmetal form ions, and it is still called an ionic compound. These metals usually only form two different ions.

From the following list, cross out those compounds that do NOT belong in the category for Type II binary ionic compounds.

AlP FeCl2 Ag2O VBr5 CoS SnF2 K3N SrF2 CuBr AuCl3 ZnO HgS

Formula and name examples for Type II binary ionic compounds:

Fe2O3 = iron(III) oxide or ferric oxide FeO = iron(II) oxide or ferrous oxide

CuS = copper(II) sulfide or cupric sulfide CuCl = copper(I) chloride or cuprous chloride

MnO2 = manganese(IV) oxide or manganic oxide MnCl2 = manganese(II) chloride or manganous chloride

What type of element is always listed first (metal or nonmetal)? ____________ second? ___________

Is the name of the first element in the compound different from the element? (yes/no)

What is the common ending for the nonmetal portion of the names? ______________

In the compound FeO, what is the charge on iron? _______

In the compound Fe2O3, what is the charge on iron? ________

What does the Roman number after the metal name represent?

Try not to look in your book and determine the rules for naming type II binary ionic compound when given the formula.

Try not to look in your book and determine the rules for writing the formula for a type II binary compound when given the name.

Name each of the type II binary ionic compounds listed in Question 1 of the Type II section in words.

Type III (Nonmetallic) Binary Compounds

Binary compounds that do not contain metals, sometimes referred to as binary molecular compounds, have covalent bonds instead of ionic bonds. A covalent bond is formed by sharing one or more pairs of electrons. The pair of electrons is shared by both atoms. For example, in forming H2, each hydrogen atom contributes one electron to the single bond.

From the following list, cross out those compounds that do NOT belong in the category for binary compounds containing only nonmetals or metalloids.

CCl4 AlCl3 CO SeF6 SiO2 SrI2 P4O10 TiO2 SeO3 IrCl ZrO2 N2O5

Formula and name examples for Type III binary molecular compounds:

CO2 = carbon dioxide H2O = dihydrogen monoxide

IF5 = iodine pentafluoride BF3 = boron trifluoride

Which element is listed first in the name?

Is the name of the first element in the compound different from the element? (yes/no)

What is the common ending for all the names? ______________

What do the prefixes (di-, mono-, penta-, tri-) in the names above mean?

Is the prefix mono- used when there is only one atom of the first element? (yes/no)

Is the prefix mono- used when there is one atom of the second element? (yes/no)

Try not to look at your book and determine the rules for naming type III binary ionic compound when given the formula.

Name each of the type III binary compounds in Question 1 of the Type III section in words.

Compounds Containing Polyatomic Ions

Polyatomic ions are ions that as a group have a set charge. Polyatomic ions are usually recognized in a formula by the grouping of more than one nonmetal elements after a metal. Your book has a table listing polyatomic ions. Use your bookâs table to fill in the following table with the appropriate names/formulas of the polyatomic ions.

Name	Formula	Name	Formula
hydronium		chlorite
phosphate			C2H3O2-1
	SO3-2		ClO4-1
	OH-1	carbonate
cyanide			NO2-1

Polyatomic ions containing oxygen (oxyanions or oxoanions) are somewhat special.

carbonate = CO3-2

Match the location of carbon on the periodic table with the two figures on the right.

What number is in the carbon location on the left figure? To what does this number refer?

What number is the carbon location on the right figure? To what does this number refer?

What element must all oxyanions contain?

What is the ending of the name of the ion determined from these tables?

Determine the formula for the following oxyanions using the figures above. Phosphate _________ Silicate ________ Bromate ________ Iodate ________ Sulfate ________ Nitrate ________

Review Table

Use your knowledge of Type I and Type II metals as well as the appropriate polyatomic name/formula to fill in the following table.

Name	Formula	Name	Formula
sodium phosphate			Cu(ClO)2
tin (II) nitrate		calcium sulfate
	KCN	ammonium nitrate
	Sr(MnO4)2		MnSO4

Acids

Acids are compounds that when dissolved in water, produce hydrogen ions (H+). Naming acids can be tricky. Use the following figure to determine how to name an acid.

FG02_27

Given that the ion formula is NO2-1, how can one determine the name of the ion, acid formula, and acid name?

Ion name:

What is the formula for nitrate?

Is NO2-1 the nitrate ion or the nitrite ion?

Acid formula:

According to the figure above, what ion must be added to create an acid? What is the charge of that ion?

How many of the hydrogen ions must be added to NO2-1 to make a neutral acid (zero charge)?

What is the chemical formula for the acid created when hydrogen ion(s) are added to NO2-1?

Acid name:

Based on your answer to Question 1 above, does the name for the NO2-1 ion end in âite or âate?

Name the acid, HNO2.

Fill in the following acid table:

Acid Formula	Acid Name	Anion Formula	Anion Name
HCl	hydrochloric acid	ClÂ¯	chloride
		ClO3Â¯
	sulfuric acid
			bromate
		PO4-3
HF
H2CO3
		C2H3O2 Â¯

roseant57

experment1 Electron Transfer (Redox) Reactions

1. Comparison of Br2/Brâ and Fe3+/Fe2+ Couples

Oxidant-reductant pairs from these couples will he tested, i.e. Br2 with the reductant Fe2+ and

Brâ with the oxidant Fe3+. A reaction will occur in only one of these cases because of the

difference in electrode potential between the two couples.

To test whether reaction has occurred, SCNâ will be used as a probe for Fe3+, with which it

gives a red complex. Note that the air rapidly oxidises Fe2+, so that a faint pink colour is often

obtained when SCNâ is added to Fe2+.

Solutions containing Fe2+ or Fe3+ tend to decompose (although in different ways); both are

stabilised by H+. The iron(II) solution provided is freshly prepared 0.1 M (NH4)2Fe(SO4)2 / 2

M H2SO4.

To 5 drops of the 0.1 M iron(II) solution in a semi-micro tube add 10 drops of 0.1 M KBr and

stir. Place 1 drop of the mixture on a white tile, and test for Fe3+ by adding 1 drop 0.1 M

KSCN. A strong red colour indicates Fe3+.

Repeat using 5 drops of the 0.1 M iron(III) solution and adding 10 drops of 0.1 M KBr.

Repeat using 5 drops of the 0.1 M iron(II) solution and adding 10 drops of 0.1 M Br2.

Repeat using 5 drops of the 0.1 M iron(III) solution and adding 10 drops of 0.1 M Br2.

2.Use of the Br2/Brâ and Fe3+/Fe2+ Couples in an Electrochemical Cell

You are supplied with a pair of carbon rods joined by copper wire. Place two 100 mL beakers

side by side.

In one put 50 mL of the iron(II) solution and 1 mL of 0.1 M KSCN solution. In the other

beaker put 25 mL of 0.2 M Br2 solution and 25 mL of 4 M HCl.

Correct any difference in level between the solutions in the two beakers by adding water to

one. Fold an 11 cm filter paper into a strip about 1 cm wide. Bend the strip in the middle,

moisten with a little 4 M HCl, then insert the strip so that it links the two solutions. Insert one

carbon rod into each solution, then allow to stand undisturbed for 30 minutes while

proceeding with Experiment 3. Carefully observe the apparatus at intervals and record any

change that occurs.

3. Comparison of the Fe3+/Fe2+ and Sn4+/Sn2+ Couples

In this experiment the additional sources of ions are:

tin(IV) ion 0.1 M SnCl4/1 M HCl

tin(II) ion 0.1 M SnCl2/1 M HCl

(Both Sn2+ and Sn4+ solutions are stabilised by H+.)

To 5 drops of the 0.1 M iron(II) solution in a semi-micro tube add 10 drops of the appropriate

0.1 M tin solution and stir. Place 1 drop of the mixture on a white tile, and test for Fe3+.

Record the results.

Similarly, test the 0.1 M iron(III) solution with an excess of the appropriate 0.1 M tin

solution.

You have now enough information to state whether the standard electrode potential of the

Fe3+/Fe2+ couple is more positive or less positive than that of the Sn4+/Sn2+ couple.

please answer following questions

1.Comparison of Br₂/Br^â and Fe³⁺/Fe²⁺ couples

Complete the following table.

Solutions	Observation	Fe³⁺ present (Yes/No)	Did a reaction occur?
Fe(II) and KBr
Fe(III) and KBr
Fe(II) and Br₂
Fe(III) and Br₂

2.Br₂/Br^â and Fe³⁺/Fe²⁺ couples in an Electrochemical Cell

a)What observations did you make over 3 minutes? Did Br₂ or Fe²⁺ react? Write the ion/electron half- equations for this reaction and the complementary reduction. Write the overall redox reaction.

b)Draw/scan/insert a circuit diagram showing the half-reactions that occur in each beaker, the direction of electron flow and the direction of flow of ions.

3.Fe³⁺/Fe²⁺ and Sn⁴⁺/Sn²⁺ couples

Complete the following table.

Solutions	Observation	Fe³⁺ present (Yes/No)	Did a reaction occur?
Fe(II) and Sn(IV)
Fe(II) and Sn(II)

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