Unit B: Electrochemical Changes
Key Concepts: The following concepts are developed in this unit and may also be addressed in other units or in the courses. The intended level and scope of treatment is defined by the outcomes.
oxidation reduction oxidizing agent
reducing agent oxidation-reduction
(redox) reaction oxidation number half-reaction disproportionation spontaneity standard reduction
potential voltaic cell electrolytic cell
electrolysis standard cell potential Faraday’s law corrosion
Specific Outcomes for Knowledge
Students will:
define oxidation and reduction operationally and theoretically
define oxidizing agent, reducing agent, oxidation number, half-reaction, disproportionation
differentiate between redox reactions and other reactions, using half-reactions and/or
oxidation numbers
identify electron transfer, oxidizing agents and reducing agents in redox reactions that
occur in everyday life, in both living systems and nonliving systems; i.e., corrosion compare the relative strengths of oxidizing and reducing agents, using empirical data
predict the spontaneity of a redox reaction, based on standard reduction potentials, and
compare their predictions to experimental results
write and balance equations for redox reactions in acidic and neutral solutions by
using half-reaction equations obtained from a standard reduction potential table
developing simple half-reaction equations from information provided about redox
changes
assigning oxidation numbers, where appropriate, to the species undergoing chemical
perform calculations to determine quantities of substances involved in redox titrations.
define anode, cathode, anion, cation, salt bridge/porous cup, electrolyte, external circuit,
power supply, voltaic cell and electrolytic cell
identify the similarities and differences between the operation of a voltaic cell and that of
an electrolytic cell
predict and write the half-reaction equation that occurs at each electrode in an
electrochemical cell
recognize that predicted reactions do not always occur; e.g., the production of chlorine gas
from the electrolysis of brine
explain that the values of standard reduction potential are all relative to 0 volts, as set for
the hydrogen electrode at standard conditions
calculate the standard cell potential for electrochemical cells
predict the spontaneity or nonspontaneity of redox reactions, based on standard cell
potential, and the relative positions of half-reaction equations on a standard reduction
potential table
calculate mass, amounts, current and time in single voltaic and electrolytic cells by applying Faraday’s law and stoichiometry.
Specific Outcomes for Science, Technology and Society (STS) (Science and Technology Emphasis)
Students will:
explain how the goal of technology is to provide solutions to practical problems describe the methods and devices used to prevent corrosion; i.e., physical coatings and
cathodic protection
explain that technological problems often require multiple solutions that involve different
designs, materials and processes and that have both intended and unintended consequences
explain that scientific knowledge may lead to the development of new technologies, and
describe science and technology applications that have developed in response to human
and environmental needs
explain that science and technology have influenced, and been influenced by, historical
development and societal needs
Specific Outcomes for Skills (Science and Technology Emphasis) Initiating and Planning
Students will:
formulate questions about observed relationships and plan investigations of questions,
ideas, problems and issues
design an experiment to determine the reactivity of various metals formulate questions about observed relationships and plan investigations of
questions,
ideas, problems and issues
design an experiment, including a labelled diagram, to test predictions regarding
spontaneity, products and the standard cell potential for reactions occurring in
electrochemical cells
describe procedures for the safe handling, storage and disposal of materials used in the
laboratory, with reference to WHMIS and consumer product labelling information
Performing and Recording
Students will:
conduct investigations into relationships among observable variables and use a broad range of tools
and techniques to gather and record data and information
select and correctly use the appropriate equipment to perform a redox titration experiment
use a standard reduction potential table as a tool when considering the spontaneity of
construct and observe electrochemical cells
Analyzing and Interpreting
Students will:
analyze data and apply mathematical and conceptual models to develop and assess possible
solutions
evaluate data from an experiment to derive a simple reduction table identify the products of electrochemical cells
compare predictions with observations of electrochemical cells identify the limitations of data collected on an electrochemical cell
explain the discrepancies between the theoretical and actual cell potential
Communication and Teamwork
Students will:
work collaboratively in addressing problems and apply the skills and conventions of science in communicating information and ideas and in assessing results
select and use appropriate numeric, symbolic, graphical and linguistic modes of representation to communicate equations for redox reactions and answers to problems
related to redox titrations
use appropriate SI notation, fundamental and derived units and significant digits to communicate answers to problems related to functioning electrochemical cells
Oxidation Number or Oxidation State
Students should be able to determine the oxidation numbers for elements in compounds including
able to determine the oxidation number for binary molecular compounds and for simple organic
molecules and carbohydrates such as sucrose and glucose. Students are not expected to be able to
assign different oxidation numbers to two atoms of the same element within a compound. For
example, each of the carbons in CH3COOH(aq) would be assigned an oxidation number of 0.
Students are expected to be able to assign oxidation numbers for simple ionic compounds, complex
ions, binary acids, and oxoacids. The terms oxidation number and oxidation state will be used
synonymously.
Balancing Electrochemical Equations
Students are expected to devise a balanced half-reaction in an acidic or neutral, but not basic,
solution. They are expected to balance chemical equations that occur in basic environments
given the species, but not to devise their own half-reactions. Students are expected to balance
disproportionation reaction equations.
Auto-oxidation versus Disproportionation
Only the term disproportionation will be used to describe a substance undergoing both an oxidation and a reduction.
Chloride Anomaly
Students are expected to recognize that predicted reactions do not always occur, for example, the
chloride anomaly that occurs during the electrolysis of solutions containing chloride ions and water
as the strongest reducing agents.
Classifications
Electrochemical cells will be defined as either voltaic cells or electrolytic cells. Line or cell notation is used to describe electrochemical cells. The convention that is used is
that the substance constituting the anode is listed at the far left, and the substance constituting
the cathode is listed at the far right.
Standard state conditions and corresponding E° values imply the use of 1.0 mol/L reagents.
Students should know that as the reaction proceeds, the voltage generated will decrease as reactants are converted to products until reaching equilibrium, at which point the battery dies.
For half-cells containing acidified solutions (such as acidified potassium permanganate and an inert
electrode), each half-cell should include all the active components in their standard state; that is,
1.0 mol/L H+(aq), 1.0 mol/L MnO4–(aq), and 1.0 mol/L Mn2+(aq). Together with an iron half-cell,
the line notation is represented as
Fe(s) | Fe2+(aq) || MnO4-(aq), Mn2+(aq), H+(aq) | Pt(s)
Use the following information to answer the next question.
Statements I Sulfur forms stable ions by gaining electrons. II Magnesium forms stable ions by losing electrons. III The oxidation number of iron changes from + 3 to + 2. IV The oxidation number of oxygen changes from – 2 to – 1.
1. The statements numbered above that refer to oxidation are A. I and III
B. I and IV C. II and III D. II and IV
Use the following information to answer the next question. Leaching technology is used in the mining and refining of copper ore. In the first step of the leaching process, concentrated aqueous sulfuric acid flows through a copper ore deposit. Solid copper(II) oxide reacts with sulfuric acid as represented by the following net ionic equation.
CuO(s) + 2H+(aq) Cu2+(aq) + H 2O(l)
The resulting solution that contains copper(II) ions is transferred to an electrolytic cell where pure copper is produced.
B. oxidation only
C. both oxidation and reduction D. neither oxidation nor reduction
3. Which of the following statements is an operational definition of the metal undergoing reduction?
A. Iron metal undergoes a formation reaction with oxygen gas. B. Magnesium metal increases in mass when heated in air.
C. Iron(III) hydroxide reacts with oxygen in the air to form ionic compounds. D. Zinc sulfide ore is roasted in the presence of oxygen gas to produce zinc metal.
Use the following information to answer the next question. 2H2S(g) + 3O2(g) 2 SO2(g) + 2H2O(g) 4. In the reaction represented by the equation above, oxygen acts as the i agent, and the oxidation number of the sulfur atom increases by ii .
The statement above is completed by the information in row
Row i ii
A. oxidizing 2
B. oxidizing 6
C. reducing 2
D. reducing 6
5. Which of the following equations represent a disproportionation reaction? A. 2 Na(s) + I2(s) 2 NaI(s)
B. 2 F2(g) + O2(g) 2 OF2(g)
C. Cl2(aq) + H2O(l) HOCl(aq) + H+(aq) + Cl–(aq)
D. 2NH3(aq) + NaOCl(aq) N2H4(aq) + NaCl(aq) + H2O(l)
Use the following information to answer the next question. The equilibrium constants for two reactions at 900 K are represented by the following equations
Equation I 2 H2S(g) + 3 O2(g) 2 SO2(g) + 2 H2O(g) + 1 036.0 kJ Kc = 377 Equation II 2 SO2(g) + O2(g) 2 SO3(g) + 197.8 kJ Kc = 13 7 . Compare the two reactions represented by the equations above in terms of redox and equilibrium. Identify one similarity or difference in terms of redox and one similarity or difference in terms of equilibrium.
Your response should include
• an explanation of the similarity and/or difference you have identified in terms of redox
• an explanation of the similarity and/or difference you have identified in terms of equilibrium
Use the following information to answer the next question. Cellular Respiration
C6H12O6(s) + 6 O2(g) 6 CO2(g) + 6 H2O(l) 8. During cellular respiration, the oxidizing agent is
A. O2(g) B. CO2(g) C. H2O(l) D. C6H12O6(s)
Metallurgical Processes I 2 Al2O3(s) 4 Al(s) + 3 O2(g)
II 2 PbO(s) + C(s) 2 Pb(s) + CO2(g) III Zn(s) + 2 HCl(aq) ZnCl2(aq) + H2(g)
IV Cu(s) + 4 HNO3(aq) Cu(NO3)2(aq) + 2 NO2(g) + 2 H2O(l) 9. The metallurgical processes in which the metal loses electrons are
A. I and II B. I and III C. II and IV D. III and IV
10. When listed in order from strongest to weakest, the oxidizing agents are A. Ra(s), Be(s), V(s), Cd(s)
B. Cd(s), V(s), Be(s), Ra(s)
C. Ra2+(aq), Be2+(aq), V2+(aq), Cd2+(aq) D. Cd2+(aq), V2+(aq), Be2+(aq), Ra2+(aq)
B. X–(aq) C. X2(aq) D. M(s)
Use the following information to answer the next question.
A student wants to store a tin(II) chloride solution in a container made of either iron, zinc, aluminium, or copper.
12. Design an experiment that will enable you to determine which type of
container—iron, zinc, aluminium, or copper—is best to use in terms of reactivity for storing a tin(II) chloride solution.
Your response should include
• an explanation and rationale to support your choice • relevant half-reaction equations
• two other criteria you would use, other than reactivity, to determine the best storage container
13. Which of the following equations represents a spontaneous redox reaction? A. Zn2+(aq) + Pb(s) Zn(s) + Pb2+(aq)
B. Sn4+(aq) + Fe(s) Sn2+(aq) + Fe2+(aq) C. Zn2+(aq) + Co(s) Zn(s) + Co2+(aq)
D. O2(g) + 2 H2O(l) + 4 Br–(aq) 2 Br2(l) + 4OH–(aq)
14. The reducing agent that can convert 1.0 mol/L Sn4+(aq) ions to Sn2+(aq) but not 1.0 mol/L Sn2+(aq) to Sn(s) is
C. Ni(s) D. Cr(s)
Use the following information to answer the next question. __OCl–(aq) + __I–(aq) + __H+(aq) _I
2(aq) + _Cl–(aq) + _H2O(l)
15. When the equation above is balanced under acidic conditions, the whole number coefficient for H+(aq) is i and the amount of electrons transferred
is ii .
The statement above is completed by the information in row
Row i ii
A. 1 1 mol
B. 1 2 mol
C. 2 1 mol
D. 2 2 mol
Use the following information to answer the next question.
Chlorine gas and aqueous sodium hyposulfite react as represented by the
following unbalanced equation.
Cl2(g) + S2O32–(aq) + H2O(l) SO42–(aq) + H+(aq) + Cl–(aq)
16. The balanced oxidation half-reaction equation is A. Cl2(g) + 2 e– 2 Cl–(aq)
B. S2O32–(aq) + H2O(l) SO42–(aq) + 2H+(aq) + 4 e–
17. Iron metal is easily oxidized to Fe2+(aq) by an acidified potassium dichromate solution during a redox titration.
a. Write the net ionic equation for this process and determine the mass of iron metal oxidized by 50.0 mL of a 0.250 mol/L acidified K2Cr2O7(aq) solution. (3 marks) b. Oxidation of iron is often an undesirable reaction in the environment. There
are several methods to prevent corrosion of iron. Describe one of these methods and explain how this method prevents iron from corroding. (2 marks)
Use the following information to answer the next two questions. A standardized 0.125 mol/L potassium dichromate solution was used to titrate 20.0 mL samples of acidified Sn2+(aq). The data is represented in the
following table.
Titration Data
Trial I II III
Final burette volume (mL) 27.2 44.5 30.1
Initial burette volume (mL) 10.1 27.2 12.9
18. The amount of potassium dichromate solution required to complete this titration is A. 8.33 x 10–4 mol
B. 6.45 x 10–3 mol C. 2.50 x 10–3 mol D. 2.15 x 10–3 mol
Use your recorded answer from Multiple Choice 18 to answer Numerical Response 19.* 19. The concentration of Sn2+(aq) in the sample used in the titration, expressed in scientific notation, is a.bc 10–dmol/L. The values of a, b, c, and d are ______, ______, ______, and
______.
(Record all four digits of your answer in the numerical-response section on the answer sheet.)
B. cations gain electrons
C. cations are attracted to the electrode D. electrons are attracted to the electrode
21. Illustrate and describe a working voltaic cell that incorporates a standard nickel half-cell and has a cell potential greater than 1.00 V.
Your response should include
• relevant balanced half-reaction equations and an E°cell calculation
• a labelled cell diagram
• evidence that a reaction has occurred in each half-cell
22. Match the numbers in the diagram above with their appropriate labels given below. Anode __________ (Record in the first column)
23. Which of the following rows identifies the type of electrochemical cell in the diagram above and describes what happens during its operation?
24. An electrolytic cell differs from a voltaic cell in that the electrolytic cell A. is spontaneous
B. consumes electricity C. has a positive E°cell value
D. has an anode and a cathode
Row Type of Cell What Happens
A. Voltaic Electrons move toward the cathode B. Voltaic I–(aq) moves toward the cathode
Use the following information to answer the next question. Statements About Electrochemical Cells
1 Oxidation occurs at the anode.
2 The oxidizing agent reacts at the cathode. 3 Cations move through the wire to the cathode. 4 Cations move through the electrolyte to the cathode. 5 Electrons move through the wire to the cathode. 6 Electrical energy is converted to chemical energy. 7 Chemical energy is converted to electrical energy.
25. The statements numbered above that apply to both electrolytic cells and voltaic cells are _____ , _____ , _____ , and _____ .
26. The reduction half-reaction that occurs during the operation of the electrochemical cell represented in the diagram above is i , and this reaction occurs at the ii .
The statement above is completed by the information in row
Row i ii
27. The half-reaction that occurs at the anode during the discharge of the nickel–cadmium cell is
A. Cd(s) + 2OH–(aq) Cd(OH)
2(s) + 2 e–
B. Cd(s) + 2OH–(aq) + 2 e– Cd(OH) 2(s)
C. NiO2(s) + 2H2O(l) + 2 e– Ni(OH)2(s) + 2OH–(aq) D. NiO2(s) + 2H2O(l) Ni(OH)2(s) + 2OH–(aq) + 2 e–
Use the following information to answer the next question. Predicted reactions do not always occur when tried in a laboratory.
28. a. Identify and describe an occurrence which may result in a predicted reaction not occurring. (2 marks)
29. For the standard reference half-cell, the reduction half-reaction equation and electrical potential are
A. H2(g) 2 H+(aq) + 2 e– E° = 0.00 V B. 2 H+(aq) + 2 e– H
2(g) E° = 0.00 V
C. 2 H2O(l) + 2 e– H2(g) + 2 OH–(aq) E° = – 0.83 V D. H2(g) + 2 OH–(aq) 2 H2O(l) + 2 e– E° = + 0.83 V
30. If the standard iodine half-cell is chosen as the reference half-cell instead of the hydrogen half-cell, then the cell potential for a silver–nickel cell is
+/– __________ V.
Use the following information to answer the next question. 2 Ag+(aq) + Zn(s) 2 Ag(s) + Zn2+(aq)
31. The cell potential for the redox reaction represented by the equation above is A. + 0.04 V
B. + 0.84 V C. + 1.56 V D. + 2.36 V
Use the following information to answer the next question. A student constructed two standard electrochemical cells using Pb2+(aq) and Ni2+(aq). In both cells a Pb(s) electrode was placed in the Pb2+(aq) solution. In the first cell a Ni(s) electrode was placed in the Ni2+(aq) solution. In the second cell an inert C(s) electrode was placed in the Ni2+(aq) solution instead of the Ni(s) electrode.
33. Which of the following statements describes what occurs in each cell? A. In both cells a power source is needed.
B. In both cells a spontaneous reaction occurs and Pb(s) is produced. C. In the first cell Ni(s) is produced, and in the second cell a power source is needed.
D. In the first cell the reaction is spontaneous, and in the second cell the reaction is nonspontaneous.
Use the following information to answer the next question. The reduction half-reaction for a Hall-Héroult electrolytic cell is represented by the following equation.
Al3+(l) + 3 e– Al(l)
34. If a current of 10.0 A is applied for 5.00 h to the Hall-Héroult electrolytic cell, then the amount of electrons transferred is
A. 5.60 mol B. 1.87 mol C. 6.22 x 10–1 mol D. 5.18 x 10–4 mol
Use the following information to answer the next question. A Hall-Héroult electrolytic cell is used to produce molten aluminium from molten aluminium oxide, as represented by the following simplified equation.
2 Al2O3(l) 4 A1(1) + 3O2(g)
at 5.55 x 103 A to produce 20.0 kg of aluminium is __________h.
Use the following information to answer the next question. A particular company manufactures plastic tape containing small pieces of magnesium. The tape is completely wrapped around iron pipes that will be buried underground.
36. Explain in chemical terms the purpose(s) of each component of the tape. Your response should include
• an explanation of the corrosion of iron
• an explanation of how the plastic tape and magnesium pieces prevent the corrosion of iron
• relevant half-reaction equations
37. From an ecological perspective, a reason why hydrogen–oxygen fuel cells should not be used to power automobiles is that
A. hydrogen fuel can be produced through the electrolysis of seawater by using the energy produced from burning fossil fuels
B. cars powered by a hydrogen–oxygen fuel cell would be up to 30 more efficient than cars powered by gasoline
D. oxygen is readily available from the atmosphere 1 D 2 D 3 D 4 B 5 C 6 B 7 DO IT 8 A 9 D 10 D 11 B 12 DO IT 13 B 14 B 15 D 16 C 17 DO IT 18 D
19 IF YOU PICKED A FROM QUESTIO 18 THE ANSWER IS 1251, B WOULD GIVE 9681, C WOULD GIVE 3751 AND D GIVES 3231
20 A 21DO IT 22 4351 23 C 24 B 25 1245 26 B 27 A 28 DO IT 29 B 30 1.06V 31 C 32 1.33V 33 D 34 B
35 10.7 OR 10.8 DEPENDING ON YOUR MATH SKILLZ AND SIG FIGS 36 DO IT
