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
Transcript
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)
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
(c) Br2 + 2KI I2 + 2 KBr
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.
(c) An aqueous chlorine solution is added to a phosphorous acid 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) + SO4
2-(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. 1.
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. 2. Add the equations 1 and 2 to create the overall reaction of iron to iron(III) performed in two steps.
Include the cell potential.
3. One Step:
Write the equation for the half-reaction of iron to iron(III) ion, including the cell potential. 4.
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. 5.
If oxygen is present in the water, it may be the oxidizing agent. Write this half-reaction including the
cell potential.
6.
In acidic solutions, hydrogen ions may act as the oxidizing agent. Write this half-reaction including the
cell potential. 7.
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. 8. 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.
9.
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.
Prediction: Use the half reactions from the previous page. Anode: 1 and 4 Cathode: 5, 6 and 7.
1. Write all the possible reactions for an Fe nail in water in an open test tube. Calculate the overall cell
potentials, determine spontaneity and describe the expected reaction observations. Verify your answers
by using the reduction table. (Note, if reaction 1 can’t occur, then 4 can’t happen either)
2. Write all the possible reactions for the spontaneous reactions for an Fe nail in acidic solution in an
open test tube. Calculate the overall cell potentials, determine spontaneity and describe the expected
reaction observations. Verify your answers by using the reduction table.
3. Write all the possible half reactions for the spontaneous reactions for an Fe nail partially covered with
Cu metal. Calculate the overall cell potentials and determine the spontaneity of the reactions.
4. Write all the possible half reactions for the spontaneous reactions for an Fe nail partially covered with
Zn metal. Calculate the overall cell potentials and determine the spontaneity of the reactions.
PROCEDURE: 1. Heat 100 mL of distilled water to the boil in a 250 mL beaker. Remove from the heat
and stir in 1 tsp of agar. Continue heating gently and stirring until the agar is dispersed. Share between two groups.
2. Add 5 drops of 0.1 mol/L potassium ferricyanide solution and 30 drops of 0.1%
phenolphthalein solution. Stir thoroughly. Allow to cool but do not allow the agar to set.
3. Place a nail covered in Cu wire and a nail covered with Zn in a Petri dish. 4. Cover the nails in the Petri dish with the agar mixture and place a labeled lid on the
dish. Store as directed. 5. Prepare four short finishing nails by washing them in soap to remove any protective
oil coating. 6. Place one small nail in each of four medium sized test tubes. To one add distilled
water. To the second add 0.1 mol/L sodium hydroxide. To the third add 0.1 mol/L hydrochloric acid. To the fourth add 0.1 mol/L sodium chloride. Set all four test tubes in a small beaker labeled with your name. Store as directed.
7. Do Not pour the extra agar solution down the drain – it will clog the drain. Scrape
out and throw in the garbage. FOLLOW-UP AND OBSERVATIONS: 8. The next lab day, make detailed diagrams on the observation sheet for the nails in the
test tube solutions. Label each observable half reaction using the numbers 1- 9 in the Introduction.
9. Test the solutions with pH paper, record. 10. Next add a drop or two of potassium ferricyanide to each test tube and note any
colour change. Record observations same as above.
11. Make detailed diagrams on the observation sheet for the nails in the test tubes. Label
each observable half reaction using the numbers 1- 9 in the Introduction.
OBSERVATIONS
Name:
Initial Diagram – Before addition of potassium ferricyanide
pH = pH = pH = pH =
Final Diagram – After addition of potassium ferricyanide
Final Diagram (1 Petri dish)
Additional Reaction
12.
Analysis: Use the half reactions and reactions from the Introduction.
Test tubes reactions Write the anode, cathode 1/2 reactions and the overall REDOX reaction including
cell potential for the Fe nail in the solution indicated.
Make sure to match the observation to the products obtained.
1. a) In H2O
b) Subsequent reactions that produced the observed product(s):
2. In HCl solution
3. In NaOH solution
4. a) In NaCl solution
b) Subsequent reactions that produced the observed product(s):
Electrolytic Cells Worksheet
1. a) Give the cathode, anode and overall equations including cell potentials to conclude what happens to
the pH of the solution near the cathode and anode during the electrolysis of KNO3? Consider all
possible reactions.
b) Write the shorthand cell notation.
2. Given the following molten systems, predict the products at each electrode. Assume inert electrodes and
sufficient voltage to cause a reaction to take place. Consider all possible rxns.
a) FeBr2
b) NiCl2
c) Na2SO4
3. Given the following 1.00 M solutions at 25°C predict the anode and cathode half cell reactions. What is
the minimum voltage required for each cell to operate?
a) LiMnO4
b) CrI3
c) Sn(NO3)2
d) Ag2SO4
Cell Stoichiometry Worksheet
1. How many coulombs, q, are required to deposit 0.587 g of Ni from a solution of Ni2+ ?
2. Three electrolysis cells are connected in series. They contain, respectively, solutions of copper (II)
nitrate, silver nitrate, and chromium (III) sulfate. If 1.00 g of copper is electrochemically deposited in the
first cell, calculate the mass of silver and chromium deposited in the other cells.
3. A constant current of 3.7 milliampere is passed through molten sodium chloride for 9.0 minutes. The
sodium produced is allowed to react with water (500 mL). What is the pH of the resulting solution?
4. Given these half-reactions and their standard reduction potentials,
2 ClO4- + 16 H+ + 14 e- Cl2 + 8 H2O Eo (ClO4
-) = + 1.47 V
S2O82- + 2 e- 2 SO4
2- Eo (S2O82-) = + 2.01 V
Calculate:
(a) Complete the REDOX reaction and calculate the Eocell.
(b) If Ca(NO3)2 (aq) is added and 2.59 g of CaSO4 is produced, calculate the pH of a 30.0 mL solution.
5. The system 2 AgI + Sn Sn2+ + 2 Ag + I- has a current of 8.46 A run through it for 1.25
minutes. Calculate the mass of silver produced.
6. Calculate the current needed to produce 5.0 mL of chlorine gas after 100. seconds at SATP for the
following reaction, if:
NiO2 + 2 Cl- + 4 H+ Cl2 + Ni2+ + 2 H2O
Review Questions for SCH 4U Electrochemistry Test
1. Balance the following REDOX reaction in acidic solution
Zn + NO3 Zn2+
+ NH4+
2. Balance the following REDOX reaction in acidic solution
MnO4 + C2O42 CO2
+ MnO2
3. Given the following reactions, generate a standard reduction potential table:
W2+ + Z Z2+ + W
X2+ + W W2+ + X
X2+ + Y no rxn
4. Describe and explain what will happen if carbon electrodes are placed in a FeCl2 solution.
Give ALL possible half reactions.
5. Use the redox spontaneity rule to predict whether the following mixtures will be spontaneous or not.
(a) Nickel metal in a solution of silver ions
(b) Chlorine gas bubbled into a bromide ion solution
(c) Copper metal in nitric acid
6. Three electrolysis cells are connected in series. They contain, respectively, solutions of zinc nitrate,
aluminum nitrate and silver nitrate. If 1.00 g of silver is deposited in the third cell what mass of
aluminum and zinc were deposited in the other cells.
7. For the cell:
Ag (s) | Ag1+ (aq) || Zn2+ (aq) | Zn (s)
a) List all possible half-reactions that will occur at the cathode, including their cell potentials.
b) List the possible half-reactions that will occur at the anode, including their cell potentials.
c) Give the full balanced REDOX reaction with the value for the cell’s Eo
d) Draw a fully labeled diagram of the electrolytic cell.
8. For the cell:
Ag (s) | S2- (aq) || HCl (aq) | Pt (s)
a) List all the possible anode reactions with their Eo values.
b) List all the possible cathode reactions with their Eo values.
c) Give the most probable reaction for the electrochemical cell and the value for the cell’s Eo
d) Draw a fully labeled diagram of the cell.
e) As this reaction proceeds, what will happen to the Eo value?
f) What would happen if HCl(aq) was added to the cathodic half-cell?
