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 SO4

2- +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- 2O + 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- 2 + 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 SO4

2- + 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- -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- a -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. Fe(CN)6 3- + Cr2O3 Fe(CN)6

4- + CrO4 2-

15. NH4NO3 N2O

16. NO2– + MnO4

– NO3– + Mn2+ (in acid solution)

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

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

19. CH4 + O2 C + H2O

20. 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(III) 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- + 12 H+ + 10 e- Cl2 + 6 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

Free Energy and Non-Standard Conditions Worksheet - AP

1. Given the following reaction:

5 S2O82- + Cl2 + 6 H2O 10 SO4

2- + 2 ClO4- + 12 H+ Eo

cell = +0.54 V

(a) ΔGo for the cell reaction.

(b) Keq for the cell reaction.

2. The system 2 AgI + Sn Sn2+ + 2 Ag + I- has a calculated Eocell = -0.015 V. What

is the value of:

(a) ΔGo for this system?

(b) Keq for this system?

3. For the following reaction, given that its standard cell potential is 0.320 V at 25oC, calculate:

NiO2 + 2 Cl- + 4 H+ Cl2 + Ni2+ + 2 H2O

(a) ΔGo for the cell reaction.

(b) Keq for the cell reaction.

4. The cell reaction:

NiO2 + 4 H+ + 2 Ag Ni2+ + 2 H2O + 2 Ag+

has Eocell = +2.48 V. What will be the cell potential at a pH of 6.00 when the concentrations of Ni2+ and

Ag+ are each 0.10 M?

5. The Eocell = + 0.135 V for the reaction

3 I2 + 5 Cr2O72- + 34 H+ 6 IO3

- + 10 Cr3+ + 17 H2O

What is Ecell if [Cr2O72-] = 0.10 M, [H1+] = 0.010 M, [IO3

1-] = 0.0010 M, and [Cr3+] = 0.00010 M?

6. A cell was set up having the following reaction.

Mg + Cd2+ Mg2+ + Cd Eocell = +1.97 V

The magnesium electrode was placed into a 1.00 M solution of MgSO4 and the cadmium electrode was

placed into a solution of unknown Cd2+concentration. The potential of the cell was measured to be

+1.88 V. What was the unknown Cd2+ concentration?

7. Suppose that a galvanic cell was set up having the net cell reaction

Zn + 2 Ag1+ Zn2+ + 2 Ag Eocell = +1.56 V

The Ag1+ and Zn2+ concentrations in their respective half cells initially are 1.00 M, and each half-cell

contains 100. mL of electrolyte solution. If this cell delivers a current at a constant rate of 0.100 A, what

will the cell potential be after 10.00 hr?

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?