Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
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Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
2
Study Guide: Things You Must Know
Vocabulary (know the definition and what it means):
 Redox reaction
 Oxidation and reduction
 Oxidizing agent and reducing agent
 Corrosion of a metal
 Oxidation number
 Electrochemistry
 Voltaic cell and electrolytic cell
 Electrode
 Anode and cathode
 Salt bridge
 Honors: electrical potential and standard cell
voltage
 Honors: standard reduction potential
 Honors: diagonal rule
 Battery, including alkaline and lead storage
battery
 Electrolysis
 Electroplating
Learning Objectives (you should know:)
 How to identify the oxidation and reduction half reactions in a redox reaction.
 How to distinguish oxidation from reduction (“Leo says Ger”).
 How to determine what is oxidized and reduced in a redox reaction.
 How to determine the oxidizing agent and reducing agent in a redox reaction.
 How to balance charge in half reactions using electrons.
 How to assign oxidation numbers to every element in a compound or ion.
 How to use oxidation number to tell if a reaction is a redox reaction.
 Which classes of reactions are always redox reactions and which are never redox reactions.
 How to balance simple redox reactions using the half reaction method.
 Honors: how to balance complex redox reactions using the half reaction method (XOHe).
 How to analyze a voltaic cell to determine:
 the oxidation and reduction half reactions
 the net balanced equation
 the direction of the flow of electrons
 the directions that cations and anions flow through the salt bridge
 the identity and sign (+ or -) of the anode and cathode (Red Cat)
 Honors: how to determine standard cell voltage of a voltaic cell using a table of standard reduction
potentials.
 How an alkaline battery and lead storage battery work
 How to analyze an electrolytic cell to determine the identity and sign of the anode and cathode and the net
reaction
 How the process of electroplating works.
Key Reference Tables
 Table J: Activity Series
 Honors: Table of Standard Reduction Potentials
Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
3
•Read Chapters 22 and 23, except all students skip “Using Oxidation Number Change” pp 663-664, and Regents
students can skip section 23.2 “Half-Cells and Cell Potentials” pp685-691.
•Lab 24: Oxidation-Reduction Reactions
•Lab 25: Electrochemistry
•Regents Tables
Table J: Activity Series
•Warm-ups and problems will be collected before you take the test. Answer all problems in the space provided.
For problems involving an equation, carry out the following steps: 1. Write the equation.
2. Substitute numbers and units. 3. Show the final answer with units. There is no credit without showing work.
Oxidation & Reduction, Oxidizing & Reducing Agents
1. Define oxidation. Define reduction.
2. Determine what is oxidized and what is reduced in each reaction. Identify the oxidizing agent and reducing
agent in each case.
a. 2Na(s) + S(s)  Na2S(s)
b. 4Al(s) + 3O2(g)  2Al2O3(s)
3. Identify these processes as either oxidation or reduction.
a. Li  Li+ + eb. 2I-  I2 + 2ec. Br2 + 2e-  2Br4. Use electronegativity values to determine which reactant is oxidized and which is reduced in each reaction.
a. H2(g) + Cl2(g)  2HCl(g)
b. S(s) + Cl2(g)  SCl2(g)
c. 2Li(s) + F2(g)  2LiF(s)
5. How does the number of electrons lost and number of electrons gained compare in every redox reaction?
6. Define corrosion, and give an example.
Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
4
Oxidation Numbers
7. Determine the oxidation number of each element in these substances.
a. S2O3
b. O2
c. Al2(SO4)3
d. Na2O2
8. Determine the oxidation number of phosphorus in each substance.
a. P4O8
b. PO4-3
c. P2O5
d. P4O6
e. PO3-2
d. S
f. H2S
9. What is the oxidation number of sulfur in each of the following:
a. H2SO4
b. SO2
c. SO3-2
10. Name the element that is oxidized, the element that is reduced, the oxidizing agent, and the reducing agent in
the following reaction:
2Fe(s) + 6HCl(aq)  2FeCl3(aq) + 3H2(g)
11. Use the changes in oxidation numbers to identify which atoms are oxidized and which are reduced in each
reaction
a. 2Na(s) + Cl2(g)  2NaCl(s)
b. 2HNO3(aq) + 6HI(aq)
 2NO(g) + 3I2(s) + 4 H2O(l)
c. 2PbSO4(s) + 2H2O(l)  Pb(s) + PbO2(s) + 2H2SO4(aq)
12. Sodium chlorite is a powerful bleaching agent used in the paper and textile industries. It is prepared by this
reaction:
4NaOH(aq) + Ca(OH)2(aq) + C(s) + 4ClO2(g)  4NaClO2(aq) + CaCO3(s) + 3H2O(l)
a. Identify the element oxidized in this reaction.
b. What is the oxidizing agent?
Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
5
Identifying Redox Reactions
13. Many decomposition, single-replacement, combination, and combustion reactions are also redox reactions.
Why is a double replacement reaction never a redox reaction?
14. Are the following redox reactions?
a. HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
b. 4Fe(s) + 3O2(g)  2Fe2O3(s)
c. 2H2O(l)  2H2(g) + O2(g)
15. Identify which of the following are oxidation-reduction reactions. If a reaction is a redox reaction, name the
element oxidized and the element reduced.
a. CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2(g) + H2O(l)
b. CuO(s) + H2(g)  Cu(s) + H2O(l)
16. Which of these unbalanced equations represent redox reactions?
a. MnO(s) + PbO2(s)  MnO4-(aq) + Pb+2(aq)
b. I2O5(s) + CO(g)  I2(s) + CO2(g)
Review of Single Replacement Reactions and Table J
17. For each pair of metals listed below, use Table J to decide which metal is more readily oxidized.
a. Fe, Cu
c. Ni, Mg
e. Pb, Zn
b. Ca, Al
d. Sn, Ag
f. Cu, Al
18. Write a balanced chemical equation, if appropriate, for what happens when a strip of copper is dipped into a
solution of iron(II) sulfate.
19. Write a balanced chemical equation, if appropriate, for what happens when an iron nail is dipped into a solution
of copper(II) sulfate?
Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
6
Balancing Redox Reactions Using the Half-Reaction Method
20. For each pair of reactants:
 predict the products using Table J.
 if a reaction occurs, write and identify the oxidation half-reaction and the reduction half reaction.
 add the half-reactions to write an overall balanced chemical equation; remember the number of electrons
in each half-reaction must cancel.
a. Al + Pb+2 
b. K+ + Sn 
c. I- + Br2 
21. Predict what will happen, if anything, when strip of aluminum is dipped into a solution of copper(II) sulfate.
Write the oxidation and reduction half-reactions for this process and the balanced equation for the overall
reaction.
22. Predict what will happen, if anything, when a strip of lead is placed in an aqueous solution of magnesium
nitrate. Write the oxidation and reduction half-reactions for this process and the balanced equation for the
overall reaction.
23. Honors Use the half-reaction method to write balanced ionic equations for the following reactions that occur in
acidic solution.
a. Sn+2(aq) + Cr2O7-2(aq)  Sn+4(aq) + Cr+3(aq)
Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
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b. I-(aq) + NO3-(aq)  I2(s) + NO(g)
d. Br2(l) + SO2(g)  Br-(aq) + SO4-2(aq)
Electrochemistry
24. How do redox reactions interconvert electrical energy and chemical energy? Describe the difference between a
voltaic cell and an electrolytic cell.
Voltaic Cells
25. For each pair of reactants:
 predict whether a reaction will occur (is the reaction spontaneous?).
 if yes, write and identify the oxidation half-reaction and the reduction half reaction.
 add the half-reactions to write an overall balanced chemical equation.
a. Ca + Pb+2 
b. Na+ + Fe 
c. Fe + Fe+3  (Hint: both reactants go to Fe+2)
d. Cl- + Br2 
Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
26. What process occurs at the anode of a voltaic cell? At the cathode?
27. Explain the function of the salt bridge in a voltaic cell.
28. Honors How was the standard reduction potential of the hydrogen electrode determined?
29. Honors What is the difference between standard cell potential and standard reduction potential?
30. A nickel half cell is connected to an iron half cell.
a. what is oxidized and what is reduced?
b. identify the anode and the cathode.
c. assign charges (positive and negative) to the electrodes.
d. write the half-reactions and the overall balanced spontaneous equation.
e. Honors calculate the standard cell potential when the half-cells are at standard conditions.
31. Honors What is the standard cell potential for the voltaic cell when the Zn/Zn+2 half cell is connected to the
Cu/Cu+2 half?
32. Determine whether these redox reactions will occur spontaneously. Honors Calculate the standard cell
potential in each case.
a. Cu(s) + 2H+(aq)  Cu+2(aq) + H2(g)
b. Mg(s) + Fe+2(aq)  Mg+2(aq) + 2Fe(s)
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Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
33. For each of the following problems
 write the half reactions taking place at the anode and cathode?
 write the net ionic reaction?
 Honors calculate Eo
a. a magnesium electrode in 1 M Mg(NO3)2 is connected to a silver electrode in 1.0 M AgNO3.
b. an iron electrode in 1 M iron(II) chloride is connected to a copper electrode in 1.0 M CuCl.
c. Honors the standard hydrogen electrode is connected to a lead electrode in 1.0 M PbSO4.
d. a zinc electrode in 1 M ZnSO4 is connected to a copper electrode in 1.0 M Cu2SO4.
34. For each of the following, write a net ionic reaction for the spontaneous reaction that takes place (if any).
a. Br2 is added to an aqueous solution of NaI.
b. Br2 is added to an aqueous solution of NaCl.
c. Magnesium is added to a strong acid (H+ ions).
d. Honors Sodium is added to water.
9
Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
10
Batteries
35. What is a battery?
36. Describe the structure of an alkaline battery. Identify the substance oxidized and the substance reduced in this
type of cell.
37. Describe the composition of the anode, cathode, and electrolyte in a fully charged lead storage battery.
Electrolytic Cells
38. What is electrolysis?
39. Describe briefly how you would electroplate a teaspoon with silver.
Review
40. Calculate the pH of solutions with the following hydrogen ion or hydroxide ion concentrations.
a. [H+] = 0.000010M
b. [OH-] = 1.0 x 10-4M
41. Identify the conjugate acid-base pairs in each equation.
NH4+(aq) + H2O(l)  NH3(aq) + H3O+(aq)
42. How many milliliters of a 4.00M KOH solution are needed to neutralize 45.0 ml of 2.50M H2SO4 solution?
43. Carbon monoxide can be removed from the air by passing it over solid diiodine pentoxide.
5CO(g) + I2O5(s)  I2(s) + 5CO2(g)
a. Identify the element being oxidized and the element being reduced.
b. How many grams of carbon monoxide can be removed from the air by 0.55 g of diiodine pentoxide (I2O5)?
Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
44. A sample of oxygen has a volume of 425 ml at 30oC. What is the new volume of the gas if the temperature is
raised to 60oC while the pressure is kept constant?
45. Concentrated nitric acid is 16M. How would you prepare 500 ml of 1.0M HNO3 starting from concentrated
nitric acid?
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Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
12
Honors: Balancing Complex Redox Reactions using the Half-reaction Method
Many redox reactions cannot be balanced by “inspection”. When inspection does not work, the following
process will give a properly balanced equation. Although the process is not difficult, all the steps must be
done correctly and in the proper order.
1. Write the unbalanced reaction.
Example:
S + HNO3  SO2 + NO + H2O
2. If the reaction is in aqueous solution, break ionic substances, strong acid and strong bases into
their ions. Do not break up oxides or polyatomic ions.
Example:
S + H+ + NO3-  SO2 + NO + H2O
3. Determine which atoms are oxidized or reduced by assigning oxidation numbers.
0
+5
+4
+2
+
Example:
S + H + NO3  SO2 + NO + H2O
Therefore S is oxidized and N is reduced.
4. Write separate half reactions for the oxidation and reduction steps.
Oxidation:
S  SO2
Reduction:
NO3-  NO
5. Balance each half reaction using “X,O,H,e” in order.
X: balance atoms other than O & H
O: balance O using water
H: balance H using H+
e: balance charge using electrons
Oxidation:
S  SO2
X:
S already balanced
O:
2H2O + S  SO2
H:
2H2O + S  SO2 + 4H+
e:
2H2O + S  SO2 + 4H+ + 4eReduction:
NO3-  NO
X:
N already balanced
O:
NO3-  NO + 2H2O
H:
4H+ + NO3-  NO + 2H2O
e:
3e + 4H+ + NO3-  NO + 2H2O
Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
6. Multiply each half reaction by an integer to make the number of electrons equal.
In this example, the lowest common denominator is 12.
Oxidation:
3x (2H2O + S  SO2 + 4H+ + 4e-)
6H2O + 3S  3SO2 + 12H+ + 12eReduction:
4x (3e- + 4H+ + NO3-  NO + 2H2O)
12e- + 16H+ + 4NO3-  4NO + 8H2O
7. “Add” the two half reactions and cancel anything that is exactly the same on both sides.
6H2O + 3S  3SO2 + 12H+ + 12e12e + 16H+ + 4NO3-  4NO + 8H2O
_____________________________________________________________________________________________________________
3S + 4H+ + 4NO3-  3SO2 + 4NO + 2H2O
8. Double check.
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Redox (Oxidation & Reduction) and Electrochemistry
Chapter 22-23 Assignment & Problem Set
14
Honors: Interpreting and Calculating Half Cell Potentials (Eo)
1. Eo’s are written as reductions.
Zn+2 + 2e-  Zn
2. Eo’s are listed for standard conditions.
1.0 M for solutions
1.0 atm. for gases
T = 25 oC
3. A redox reaction is spontaneous in the forward direction if Ecell is positive.
Zn + Cu+2  Zn+2 + Cu
Ecell = 1.10 V (spontaneous)
Zn+2 + Cu  Zn + Cu+2
Ecell = -1.10 V (nonspontaneous)
4. All half-reactions are reversible. Any electrode can act as the anode or cathode depending on what it is
coupled with.
5. Table N (attached) is an upside-down version of Table J, containing standard reduction potentials for
half reactions. On Table N, the higher element is reduced and the lower element is oxidized, just
the reverse of Table J.
6. Changing the stoichiometric coefficients of a half-cell does NOT change Eo.
I2 + 2e-  2I- Eo = 0.53 V
2I2 + 4e-  4I- Eo = 0.53 V
7. Reversing the direction of a reaction changes the sign of the half-cell voltage.
I2 + 2e-  2I- 0.53 V
2I-  I2 + 2e- -0.53 V
8. These rules apply whether the reactants are separated in an electrochemical cell or mixed together.
Eocell = Eored + Eoox
9.
Voltage
measured
between two
half cells.
Std. reduction
potential of the
reduction half
cell.
Std. reduction
potential of the
oxidation half cell
with sign switched.