Balancing Redox Equations Exercise

ELECTROCHEMISTRY

57. REDOX & OXIDATION NUMBERS

An overview of oxidation and reduction processes and the assignment of oxidation numbers.

Class Work / Homework Text: p. 650 - 663

EXERCISE – Assigning Oxidation Numbers (6)

p. 649 Are You Ready? 3 - 8 p. 653 Practice 2

p. 656 Practice 8 – 11 p. 659 Practice 12 - 17

p. 662 Practice 18 – 20 p. 663 Section 9.1 Questions 1 - 10

Read: p. 664 – 673

58. BALANCING REDOX EQUATIONS

The balancing of redox equations using the half-cell and the oxidation numbers methods.

Class Work / Homework Text: p. 664 - 673

EXERCISE – Balancing Redox Equations (6)

p. 668 Practice 2 – 4 p. 673 Practice 6 - 7

p. 673 Section 9.2 Questions 3, 4, 6c

Read: p. 674 - 684

59. PREDICTING REDOX REACTIONS

The development of the Activity Series and its use to predict redox reactions.

Class Work / Homework Text: p. 674 - 684

EXERCISE – Predicting Redox Reactions (2)

p. 676 Practice 2 – 9 p. 678 Practice 10, 11, 13

p. 679 Practice 14 - 17, 19 p. 681-2 Practice 23 – 26

p. 683 Section 9.3 Questions 5 - 12

Read: p. 695 - 700

60. INTRODUCTION TO ELECTROCHEMICAL (GALVANIC) CELLS

The structure and function of galvanic cells, the methods of representing them, half cell potentials and the calculation of cell potentials both standard and non standard.

Class Work / Homework Text: p. 695 - 700

EXERCISE – Electrochemical (Galvanic or Voltaic) Cells Worksheet (2)

p. 700 Practice 1 - 9

Read: p. 700 - 709

61. REACTIONS WITH METALS

The REDOX reactions with standard hydrogen half cell, potentials and the calculation of cell potentials both standard and non standard.

Class Work / Homework Text: p. 700 - 709

Lab Preparation – Complete the INTRODUCTION questions and read the PROCEDURE for CORROSION of IRON

p. 708 Practice 10 – 16 p. 708 Section 9.5 Questions 1 - 8

Read: p. 685 – 694 and p. 710 - 714

62. CORROSION - APPLIED Part 1

63. CORROSION - APPLIED Part 2

A study of the nature of iron corrosion and the factors affecting corrosion.

Class Work / Homework Text: p. 685 – 694 and p. 710 - 714

Day 1:

Lab – Complete the PREDICTION questions for CORROSION of IRON

Day 2:

Lab – Complete the OBSERVATIONS and ANALYSIS for CORROSION of IRON

p. 687 Practice 1 - 4 p. 690 Practice 9, 10

p. 693 Practice 15 – 20 p. 694 Section 9.4 Questions 1 – 5

p. 714 Section 9.6 Questions 1 - 8, 10

Read: p. 728 - 746

64. ELECTROLYTIC CELLS AND ELECTROPLATING

The structure, design, chemistry and use of electrolytic cells. A demonstration of the process of electroplating.

Class Work / Homework Text: p. 728 - 746

EXERCISE – Electrolytic Cells Worksheet (2)

p. 735 Practice 1 - 4 p. 741 Practice 1 - 3

p. 744 Practice 8 - 10, 12 p. 776 Section 10.1 Questions 1 – 9

p. 746 Section 10.2 Questions 1, 6 - 8

Text Error

Pg 730 Section 10.1 ∆E°cell = -0.63V NOT –0.89V

Read: p. 747 - 753

65. STOICHIOMETRY OF CELL REACTIONS

Faraday’s Law and the stoichiometry of electrolytic reactions.

Class Work / Homework Text: p. 747 - 753

EXERCISE – Cell Stoichiometry Worksheet (2)

EXERCISE – Free Energy and Non-Standard Conditions Worksheet – AP (2)

p. 748 Practice 1 - 4 p. 749 Practice 5 - 7

p. 751 Practice 8 – 12 p. 752 Practice 13
p. 757 Section 10.3 Questions 1 - 7

66. REVIEW

Class Work / Homework

EXERCISE – Review Questions of SCH 4U Electrochemistry Test (4)

67. LAB TEST/APPLICATION

68. TEST

Standard Reduction Potentials

Half-Reaction E° (volts) Half-Reactions E° (volts)

F2 + 2 e- « 2 F-1 +2.87 Fe3+ + 3 e- « Fe -0.04

S2O82- + 2 e- « 2 SO42- +2.01 Pb2+ + 2 e- « Pb -0.13

Co3+ + e-1 « Co2+ +1.81 Sn2+ + 2 e- « Sn -0.14

Pb4+ + 2 e-1 « Pb2+ +1.80 AgI + e- « Ag + I-1 -0.15

H2O2 + 2 H+ + 2 e- « 2 H2O +1.77 Ni2+ + 2 e- « Ni -0.26

Au+ + e- « Au +1.69 Co2+ + 2 e- « Co -0.28

PbO2 + SO42- + 4H+ + 2e- « PbSO4 + 2 H2O +1.69 H3PO4 + 2 H+ + 2 e- « H3PO4 + H2O -0.28

MnO41- + 8 H+ + 5 e- « Mn2+ + 4 H2O +1.51 Tl+ + e- « Tl -0.34

Au3+ + 3 e- « Au +1.50 PbSO4 + 2 e- « Pb + SO42- -0.36

Ce4+ + e- « Ce3+ +1.44 Se + 2 H+ + 2 e- « H2Se -0.40

ClO41- + 8 H+ + 8 e- « Cl1- + 4 H2O +1.39 Cd2+ + 2 e- « Cd -0.40

Cl2 + 2 e- « 2 Cl- +1.36 Cr3+ + e- « Cr2+ -0.41

2 HNO2 + 4 H+ + 4 e- « N2O + 3 H2O +1.30 Fe2+ + 2 e- « Fe -0.45

Cr2O72- + 14 H+ + 6 e- « 2 Cr3+ + 7 H2O +1.23 S + 2 e- « S2- -0.48

O2 +4 H+ + 4 e- « 2 H2O +1.23 Ga3+ + 3 e- « Ga -0.53

MnO2 + 4 H+ + 2 e- « Mn2+ + 2 H2O +1.22 Ag2S + 2 e- « 2 Ag + S2- -0.69

2 IO31- + 12 H+ + 10 e- « I2 + 6 H2O +1.20 Cr3+ + 3 e- « Cr -0.74

Br2 + 2 e- « 2 Br-1 +1.07 Zn2+ + 2 e- « Zn -0.76

AuCl41- + 3 e- « Au + 4 Cl- +1.00 Te + 2 H+ + 2 e- « H2Te -0.79

Hg2+ + 2 e- « Hg +0.85 2 H2O + 2 e- « 2 OH1- + H2 -0.83

ClO1- + H2O + 2 e- « Cl1- + 2 OH-1 +0.84 Cr2+ + 2 e- « Cr -0.91

Ag+ + e- « Ag +0.80 Se + 2 e- « Se2- -0.92

NO31- + 2 H+ + e- « NO2 + H2O +0.80 SO42- + H2O + 2 e- « SO32- + 2 OH-1 -0.93

Hg22+ + 2 e- « 2 Hg +0.79 Te + 2 e- « Te2- -1.14

Fe3+ + e- « Fe2+ +0.77 Mn2+ + 2 e- « Mn -1.18

O2 + 2 H+ +2 e- « H2O2 +0.70 V2+ + 2 e- « V -1.19

MnO41- + 2 H2O + 3 e- « MnO2 + 4 OH-1 +0.60 Al3+ + 3 e- « Al -1.66

I2 + 2 e- « 2 I- +0.54 Ti2+ + 2 e- « Ti -1.75

Cu+ + e- « Cu +0.52 Be2+ + 2 e- « Be -1.85

O2 + 2 H2O + 4 e- « 4 OH- +0.40 Mg2+ + 2 e- « Mg -2.37

Cu2+ + 2 e- « Cu +0.34 Ce3+ + 3 e- « Ce -2.48

SO42- + 4 H+ +2 e- « SO2 + 2 H2O +0.18 Na+ + e- « Na -2.71

SO42- + 4 H+ +2 e- « H2SO3 + H2O +0.17 Ca2+ + 2 e- « Ca -2.87

Sn4+ + 2 e- « Sn2+ +0.15 Ba2+ + 2 e- « Ba -2.91

Cu2+ + e- « Cu+ +0.15 Cs+ + e- « Cs -2.92

S + 2 H+ + 2 e- « H2S +0.14 Ra2+ + 2 e- « Ra -2.92

AgBr + e- « Ag + Br-1 +0.07 K+ + e- « K -2.92

2 H+ + 2 e- « H2 +0.00 Li+ + e- « Li -3.00

Assigning Oxidation Numbers and Balancing Redox Equations

1. Ag + NO3 - « Ag1+ + NO

2. N2H4 + H2O2 « N2 + H2O

3. CO + Fe2O3 « FeO + CO2

4. NO3 - + CO « CO2 + NO2

5. H2 + Fe3O4 « Fe + H2O

6. H2C2O4 + MnO4 - « CO2 + MnO

7. Zn + NO3 - « Zn2+ + NO

8. C2N2 « CN - + CNO –

9. ClO2 + SbO2 - « ClO2 - + Sb(OH)6 –

10. Cr2O7 2- + I - « Cr3+ + I2

11. Fe3O4 + H2O2 « Fe3+ + H2O

12. MnO4 - + NH3 « MnO2 + NO3 –

13. CN - + CrO4 2- « CNO - + Cr(OH)3

14. NH4NO3 « N2O

15. NO2– + MnO4– « NO3– + Mn2+ (in acid solution)

16. I- + MnO4- « I2 + MnO2 (in basic solution)

17. Cl2 + S2O32- « Cl- + SO42- (in acidic solution)

18. Br2 « Br- + BrO3- (in basic solution)

Predicting REDOX Reactions

Building a REDOX Table

1. The following reactions were performed. Construct a table of relative strengths of oxidizing and reducing agents written as reductions and with the SOA to WOA.

Zn + Co2+ « Zn2+ + Co

Mg2+ + Zn « no rxn

2. In a school laboratory four metals were combined with each of four solutions. Construct a table of relative strengths of oxidizing and reducing agents written as reductions and with the SOA to WOA.

Be + Cd2+ « Be2+ + Cd

Cd + 2 H+ « Cd2+ + H2

Ca2+ + Be « no rxn

Cu + 2 H+ « no rxn

3. Write and rank the two half reaction equations for each of the following reactions:

(a) Co + Cu(NO3)2 « Cu + Co(NO3)2

(b) Cd + Zn(NO3)2 « Zn + Cd(NO3)2

4. Prepare a REDOX table of half-reactions showing the relative strengths of oxidizing and reducing agent for the following:

Al3+ / Tl+ / Ga2+ / In3+
Al / X / √ / √ / √
Tl / X / X / X / X
Ga / X / √ / X / √
In / X / √ / X / X

Prediction REDOX Reaction in Solution

1. List all the entities initially present in the following mixtures and identify all possible oxidizing and reducing agents. Write the resulting REDOX reaction (or no rxn).

(a) A lead strip is placed in a copper (II) sulfate solution.

(b) A potassium dichromate solution is added to an acidic iron (II) nitrate solution.

(d) A potassium permanganate solution is mixed with an acidified tin (II) chloride solution.

Electrochemical (Galvanic or Voltaic) Cells Worksheet

1. a) Determine the anode, cathode and calculate the standard cell potential produced by a galvanic cell consisting of a Ni electrode in contact with a solution of Ni2+ ions and a Ag electrode in contact with a solution of Ag1+ ions.

b) Write the shorthand cell notation.

2. a) Determine the anode, cathode and calculate the voltage produced by a galvanic cell consisting of an Fe electrode in contact with a solution of Fe2+ ions and a Al electrode in contact with a solution of Al3+ ions.

b) Write the shorthand cell notation.

3. a) Determine the anode, cathode and calculate standard cell potential produced by a galvanic cell consisting of a C electrode in contact with an acidic solution of ClO4- ions and a Cu electrode in contact with a solution of Cu2+ ions. Which is anode and which is the cathode?

b) Write the shorthand cell notation.

4. An electrochemical cell is constructed using electrodes based on the following half reactions:

Pb2+ (aq) + 2e- « Pb(s) Au3+(aq) + 3e « Au(s)

a) Which is the anode and which is the cathode in this cell?

b) What is the standard cell potential?

5. Use complete half-reactions and potentials to predict whether the following reactions are spontaneous or non-spontaneous in aqueous solutions. If the cell is spontaneous, write the cell shorthard notation.

a) Ca2+(aq) + 2 I-(aq) « Ca(s) + I2(aq)

b) 2 H2S(g) + O2(g) « 2 H2O(l) + 2 S(s)

c) SO2(g) + MnO2(s) « Mn2+(aq) + SO42-(aq)

d) 2 H+(aq) + 2 Br-(aq) « H2(g) + Br2(aq)

e) Ce4+(aq) + Fe2+(aq) « Ce3+(aq) + Fe3+(aq)

f) Cr2+(aq) + Cu2+(aq) « Cr3+(aq) + Cu+(aq)

CORROSION OF IRON

INTRODUCTION:

When the surface of iron is wet it undergoes oxidation. In this lab you will investigate whether the anode reaction is a two-step process of iron to the iron(II) ion followed by the formation of the iron(II) ion. It can also proceed as a one-step process of iron to iron(III) ion.

Anode (oxidation)

Two Step:

Write the equation for the half-reaction of iron to iron(II) ion, including the cell potential.

The iron(II) ion may further oxidize to the iron(III) ion.

Write an equation for the half-reaction in which oxidation of iron(II) ion to iron(III) ion occurs, including the cell potential.

Add the equations 1 and 2 to create the overall reaction of iron to iron(III) performed in two steps. Include the cell potential.

One Step:

Write the equation for the half-reaction of iron to iron(III) ion, including the cell potential.

Cathode (reduction)

Cathodic (reduction) points occur where the iron is in contact with a metal which has a higher reduction potential, or where the water has a high concentration of some oxidizing agent.

Write the half-reaction of neutral water, acting as an oxidizing agent. Include the cell potential.

If oxygen is present in the water, it may be the oxidizing agent. Write this half-reaction including the cell potential.

In acidic solutions, hydrogen ions may act as the oxidizing agent. Write this half-reaction including the cell potential.

Complete Redox reaction

In water, OH- ions produced in the cathode (reduction) reaction, will combine with the iron ions from the anode (oxidation) reaction and precipitate from the solution. The iron(II) hydroxide is not seen, but the ferric hydroxide precipitates. Write an equation for the precipitation of iron(III) hydroxide from a solution of iron(III) ion.

If the iron (III) hydroxide is dehydrated (lose water), it forms iron(III) oxide or rust.

Write an equation for the dehydration of iron(III) hydroxide.

Indicators

Potassium ferricyanide, K3Fe(CN)6, can be used to detect the presence of iron(II). (You will have to research the answer to this one). Give both the chemical equation and the 2 colours in solution, not solid.

10.

Phenolphthalein, a common indicator test (often used in titrations) that could be used to identify the products of the cathodic half-reaction. Give both the equation and the 2 colours.

11.

Overall

If the solution is acidic, the OH- ions will be neutralized, consider acid/base reactions. If the solution is basic, the production of the OH- is suppressed, consider common ion and equilibrium situations. But if the solution is neutral, the OH- ions produced in this half-reaction migrate toward the anodic points. Fe2+ and Fe2+ ions migrate from the anode to the cathode. Somewhere in-between the iron ions meet hydroxide ions and form the hydroxides, dehydrate and form rust. The water on the surface of the iron acts as an electrolyte, transporting ions between the anodic and cathodic points. Any dissolved salts present in the water will aid in this charge transfer and so accelerate the corrosion.