Chapter 20: Electrochemistry. Tentative Schedule
Date
One
Topics
Review tests, introduce red-ox,
identify oxidizing and reducing
agents
Problems
Oxidation numbers /
oxidizing and reducing
agents worksheet
Video(s) Due
1
Two
Sections 20.1-2
9, 10, 17a,c,e,g, 19a,b,e,f
2
Three
Section 20.3-4
Voltaic (Galvanic) Cells
Cell worksheet, 23, 25, 26,
27, 29
3,4
Cell emf and standard reduction
potentials
Four
Quiz 1 on Galvanic Cell
Description
Section 20.4
28,31,33,35, 36
5,6
Five
Quiz 2 on Galvanic Cell
Description
Oxidizing and Reducing Agents
39,41, 43,93
7
Six
Quiz 3 on Galvanic Cell
Description
Spontaneity and K
47, 49, 51, 53
8, 9
Seven
Nernst Equation Section 20.5-6
55, 57,59,61, 92 and
10, 11
THQ (due date for THQ
t.b.a.)
62,79,99,100
12
83, 84,85,86, *105 a,b
13
Eight
Section 20.6-8
Review Nernst and peruse
Batteries and Corrosion
Possible Al-air battery lab
Nine
Ten
Section 20.9
Electrolysis
Electrolysis again
Eleven
Lab: Electrochemistry
Twelve
Electrochemistry Review
Thirteen
Test on Electrochemistry
Electrolytic Cell
Worksheet
Practice Tests and Take
Home Tests
Written and Multiple
Choice 
Worksheet: Oxidation numbers / oxidizing and reducing agents worksheet
1)
Review pages 128-129 in your textbook on “Assigning Oxidation Numbers.”
Using these rules, identify the oxidation state of the BOLD element in each of
the following:
a) NaH
b) Cr2O72c) IF4
d) SO2
e) SO3
f) COCl2
g) MnO4h) HBrO
i) C2O42-
2)
For each reaction, very neatly:
i) identify the oxidation state of each element
ii) the element oxidized and the element reduced in each reaction
iii) the oxidizing agent and the reducing agent in each reaction (for help with
this, and examples, see page 779 in your text)
b) 2Fe(NO3)2(aq) +2Al(s)
c) Cl2(aq) + 2NaI(aq)
3Fe(s) + 2Al(NO3)3(aq)
I2(aq) + 2NaCl(aq)
d) PbS(s) + 4H2O2(aq)
PbSO4(s) + 4H2O(l)
e) Fe2O3(s) + 3CO(g)
2Fe(s) + 3CO2(g)
2--
--
Cr3+(aq) + Cl2(g)
--
CNO--(aq) + MnO2(s)
f) Cr2O7 + Cl (aq)
--
g) CN + MnO4 (aq)
Sample Problems Chapter 20
OIL RIG or LEO GER
Oxidizing agent or oxidant
Reducing agent or reductant
So, how do you know if a reaction is redox or not?
What are the oxidation states of the atoms in each substance: CN-1, HNO3, S8, SCl2, SO42-
Well, what kinds of reactions are redox?
Presenting: The Happy Life of an Electron in a Redox Reaction
Mg + CuSO4 
Mg + Cu+2 
Cu + MgSO4 
Cu + Mg+2 
Can the electrons that are transferred by used for anything?
The nickel-cadmium (nicad) battery, a rechargeable “dry cell” used in battery operated devices, uses the
following redox reaction to generate electricity:
Cd(s)+NiO2(s)+2H2O(l)
Cd(OH)2(s)+Ni(OH)2(s)
Identify the substances that are oxidized and reduced, and indicate which are oxidizing agents and which
are reducing agents.
Lots and Lots more notes must be taken on your own pieces of paper for the next bunch of videocasts.
Enjoy!
Identify the oxidizing and reducing agents in the redox reaction:
2H2O(l)+Al(s)+MnO4--(aq)
Al(OH)4--(aq)+MnO2(s)
Complete and balance this equation by the method of half reactions:
Cr2O72--(aq)+Cl --(aq)
Cr3+(aq)+Cl2(g) (acidic solution)
Complete and balance this equation by the method of half reactions:
Cu(s)+NO3--(aq)
Cu2+(aq)+NO2(g) (acidic solution)
Complete and balance this equation by the method of half reactions:
CN--(aq)+MnO4--(aq)
CNO--(aq)+MnO2(s) (basic solution)
Electrode A
Electrode B
20.3 Why do anions in a salt bridge migrate toward the anode?
The redox reaction:
Cr2O72--(aq)+14H+(aq)+6I --(aq)
2Cr3+(aq)+3I2(s)+7H2O(l)
is spontaneous. A solution containing K2Cr2O7 and H2SO4 is poured into one beaker, and a solution of KI is
poured into another. A salt bridge is used to join the beakers. A metallic conductor that will not react with
either solution (such as platinum foil) is suspended in each solution and the two conductors are connected
with wires through a voltmeter or some other device to detect and electric current. The resultant voltaic cell
generates and electric current. Indicate:
a) the reaction at the
a. anode
b. cathode
b) the direction of electron migration
c) the direction of ion migration
d) the signs of the electrodes
20.4 Cell EMF under standard conditions
For the Zn=Cu+2 voltaic cell shown above, we have
۫ cell=1.10V
Zn(s)+Cu2+(aq,1M)
Zn2+(aq,1M)+Cu(s) E◌˚
2+
Given that the standard reduction potential of Zn to Zn(s) is -0.76V, calculate the E˚red for the reduction of
Cu2+ to Cu(s).
Line notation is sometimes used to identify each cell in an electrochemical reaction to list its components as
follows:
anode anode solution
cathode solution cathode
A double vertical line represents a salt bridge or a porous barrier. A single vertical line represents a change
in phase, such as from solid to a solution.
Zn Zn2+ H+ H2
If both aqueous substances are 1.0M. Indicate:
e) the reaction at the
a. anode
b. cathode
f) the direction of electron migration
g) the direction of ion migration
h) the signs of the electrodes
i) calculate the cell emf
Calculate the standard emf of the voltaic cell based on the reaction:
Cr2O72--(aq)+14H+(aq)+6I --(aq)
2Cr3+(aq)+3I2(s)+7H2O(l)
A voltaic cell is based on:
Fe Fe+2 Ag+ Ag
a) What reaction occurs at the anode?
b) What reaction occurs at the cathode?
c) Draw this cell and identify the flow of electrons and migration of ions
d) What is the standard cell potential?
Using the standard reduction potential tables, rank the following ions in order of increasing strength as
oxidizing agents:
NO3--(aq) , Ag+(aq) , Cr2O72--(aq)
Using the standard reduction potential tables, rank the following ions in order of increasing strength as
reducing agents:
I-(aq), Fe(s), Al(s)
20.5 Free Energy and Redox
Using standard reduction potentials, determine whether the reactions are spontaneous under standard
conditions:
+
+
a) Cu(s) + 2H (aq)+>Cu2 (aq) +H2(g)
--
b) Cl2(g)+2I (aq)
2Cl --(aq)+I2(s)
Which is the stronger reducing agent, Hg(l) or Pb(s)?
Calculate the free energy change and the equilibrium constant, K, at 298K for:
4Ag(s)+O2(g)+4H+(aq)
4Ag+(aq)+2H2O(l)
Find E˚, ΔG˚, and K when the reaction is written:
2Ag(s)+12O2(g)+2H+(aq)
2Ag+(aq)+H2O(l)
Calculate the emf at 298K generated by:
Cr2O72--(aq)+14H+(aq)+6I --(aq)
2Cr3+(aq)+3I2(s)+7H2O(l)
When the concentrations of each specie are (in order of reaction):
2.0M, 1.0M, 1.0M, 1.0x10-5M, (I2(s) and H2O(l) do, obviously, not have concentrations)
If the voltage of the above cell is 0.45V at 298K when [Zn+2]=1.0M and PH2=1.0atm, what is the
concentration of H+(aq)?
What is the pH of the solution in the cathode compartment of the above cell when when [Zn+2]=0.10M and
PH2=1.0atm and the cell emf is 0.542V? You will need extra paper of your own to write this down.
You need paper of your own to write notes.
Electrolysis Sample Problems
1) Calculate the number of grams of aluminum produced in 1.00 h by the electrolysis of molten AlCl3
if the electrical current is 10.0A.
2) How many seconds would be required to produce 48.6g of Mg from MgCl2 if the current is
100.0A?
3) Calculate the number of kilowatt hours of electricity required to produce 1.0 x 103 kg of aluminum
by electrolysis of Al+3 if the applied voltage is 4.50V.
4) The compound NaI dissolves in pure water according to the equation
NaI(s) → Na+(aq) + I-(aq).
Some of the information in the table of standard reduction
potentials given below may be useful in answering the questions that follow.
(d) An electric current is applied to a 1.0 M NaI solution.
(i) Write the balanced oxidation half-reaction for the reaction that takes place.
(ii) Write the balanced reduction half-reaction for the reaction that takes place.
(iii) Which reaction takes place at the anode, the oxidation reaction or the
reduction reaction?
(iv) All electrolysis reactions have the same sign for ΔG°. Is the sign positive
or negative? Justify your answer.
Electrochemical Cell Quiz
Electrode B
Electrode A
Given the following information, answer the following about the electrochemical cell above:
Electrode A is made of lead and the solution is lead (II) nitrate
Electrode B is made of manganese and the solution it is in is manganese (II) nitrate
1) Which is the most easily reduced metal?
2) What is the balanced equation showing the spontaneous reaction that occurs?
3) What is the maximum emf the cell can produce if everything exists at standard
conditions?
4) What is the direction of flow of electrons in the wire?
5) What is the direction of the positive ion flow in the salt bridge?
6) Which electrode is increasing in size?
7) Which electrode is decreasing in size?
8) What is happening to the concentration of the ions in electrode A’s solution?
9) Identify the
i)
anode
ii)
cathode
iii)
positive electrode
iv)
negative electrode
If electrode “A” were replaced with a(n) S.H.E., what would be
10) emf of the cell
11) direction of electron flow
Electrolytic Cell Worksheet
1) Molten NiBr2(l) is electrolyzed in an electrolytic cell.
a) write the balanced chemical equation for the reaction.
b) write the half reaction that occurs at:
i) the anode
ii) the cathode
c) if a solution of aqueous NiBr2(aq) were electrolyzed in an electrolytic cell
i) write the half reaction that occurs at the anode
ii) write the half reaction that occurs at the cathode
2) Write and equation that relates Coulombs, amperes and time.
3) How many Coulombs of charge are in one mole of electrons? What is the special name
given to this amount of charge?
4) i) Write the half reaction for the formation of magnesium metoal upon electrolysis of
molten MgCl2.
ii) Calculate the mass of Mg formed upon passeage of a current of 60.0A for a period of
4.00x103 s.
iii) How many seconds would be required to produce 50.0g of Mg from MgCl2 if the
current is 100.0A?
5) In an electrolytic cell, a current of 0.250 ampere is passed through a solution of a
chloride of iron, producing Fe(s) and Cl2(g).
(a) Write the equation for the half-reaction that occurs at the anode.
(b) When the cell operates for 2.00 hours, 0.521 gram of iron is deposited at one
electrode. Determine the formula of the chloride of iron in the original solution.
(c) Write the balanced equation for the overall reaction that occurs in the cell.
(d) How many liters of Cl2(g), measured at 25˚C and 750 mm Hg, are produced when the
cell operates as described in part (b) ?
(e) Calculate the current that would produce chlorine gas from the solution at a rate of
3.00 grams per hour.
Bonus problem (show all work and units):
20.107 on p.897.
Take Home Quiz
1) A galvanic cell is constructed using a chromium electrode in a 1.00-molar solution of Cr(NO3)3 and a
copper electrode in a 1.00-molar solution of Cu(NO3)2. Both solutions are at 25˚C.
(a) Write a balanced net ionic equation for the spontaneous reaction that occurs as the cell operates.
Identify the oxidizing agent and the reducing agent.
(b) A partial diagram of the cell is shown below.
Cr
1.0 M Cr(NO 3)3
Cu
1.0 M Cu(NO 3) 3
(i) Which metal is the cathode?
(ii) What additional component is necessary to make the cell operate?
(iii) What function does the component in (ii) serve?
(c) How does the potential of this cell change if the concentration of Cr(NO3)3 is changed to 3.00-molar at
25˚C? Explain.
2) Answer the following questions regarding the electrochemical cell shown below.
(a)
(b)
(c)
(d)
(e)
Write the balanced net-ionic equation for the spontaneous reaction that occurs as the cell operates, and
determine the cell voltage.
In which direction do anions flow in the salt bridge as the cell operates? Justify your answer.
If 10.0 mL of 3.0-molar AgNO3 solution is added to the half-cell on the right, what will happen to the
cell voltage? Explain.
If 1.0 gram of solid NaCl is added to each half-cell, what will happen to the cell voltage? Explain.
If 20.0 mL of distilled water is added to both half-cells, the cell voltage decreases. Explain.
Electrochemistry Multiple Choice
(Taken from Barron’s 3rd Ed and Cliffs Notes and Arco 2nd Ed)
1) Balance the following half-reaction in acidic solution:
NO3-  NH4+
When balanced with the smallest whole-number coefficients possible, the sum of all the coefficients is
a) 13 b) 26 c) 15 d) 23 e) 21
2) Which of the following compounds includes and element with an oxidation number of +5?
a) ClO4- b) MnO4- c) NO2- d) SO32- e) NO33) Which of the following metals does NOT react with water to produce hydrogen?
a) Zn b) Li c) Ca d) Na e) K
4) The standard voltage, E˚, is 0.68V for the cell: In| In+3 || Cu+2| Cu
Find the standard reduction potential for the reaction: In+3 + 3e-  In
a) 0.34V b) –0.34V c) 1.02V d) –1.02V e) 0.00V
5) Given the following information, which of the statements is true?
Cu2+ (aq) + e-  Cu+(aq)
E˚red= 0.34V
2H+(aq) + 2e-  H2(g)
E˚red= 0.0V
Fe+2(aq) + 2e-  Fe(s)
E˚red= -0.44V
Ni(s)  Ni+2(aq) + 2eE˚red= 0.25V
a)
b)
c)
d)
e)
Cu+2(aq) is the strongest oxidizing agent
Cu+2(aq) is the weakest oxidizing agent
Ni(s) is the strongest oxidizing agent
Fe(s) would be the weakest reducing agent
H+(aq) would be the strongest oxidizing agent
6) Look at the figure below.
Which of the following is FALSE.
a) The magnesium electrode is being oxidized to Mg+2
b) The aluminum electrode is the cathode
c) The magnesium electrode is the anode
d) The aluminum ions are being reduced
e) The positive ions will flow from the Al compartment to the Mg compartment
7) Given that
Zn+2(aq) + 2e-  Zn(s)
E˚red= -0.76V
Cr3+ (aq) + 3e-  Cr(s)
E˚red= -0.74V
Calculate the equilibrium constant K for the following balanced reaction:
3Zn(s) + 2Cr+3(aq)  3Zn+2(aq) + 2Cr(s)
a)
b)
c)
d)
e)
e-0.02
e0.02
e4.7
e8.0
cannot be determined from the information provided
8) An electric current is applied to an aqueous solution of FeCl2 and ZnCl2. Which reaction occurs at the
cathode?
a) Fe2+(aq) + 2e-(aq)  Fe(s)
E˚red=-0.44V
b) Fe(s)  Fe2+ (aq) + 2eE˚ox=0.44V
c) Zn+2(aq) + 2e-(aq)  Zn(s)
E˚red=-0.76V
d) Zn(s)  Zn2+ (aq) + 2eE˚ox=0.76V
e) 2H2O(l)  O2(g) + 4H+(aq) + 4e- E˚ox=-1.23V
9) What is the number of Faradays (moles of electrons) required to produce 9.0g of aluminum by the
electrolysis of molten aluminum oxide, Al2O3?
a) 9 F b) 4F c) 3F d) 2F e) 1F
10) A voltaic cell is set up using the system
Fe|Fe+2 || Ag+|Ag
The cathode reaction produces:
a) Fe b) Fe2+ c) Ag+ d) Ag e) H2
11) A voltaic cell has an E˚cell of 1.56V and the overall reaction:
Zn(s) + 2Ag+(aq)  Zn2+(aq) + 2Ag(s)
The [Zn2 ] is 0.00010M and the [Ag+] is 0.10M. What is the voltage, E, of this cell?
a) 1.65V b) 1.62V c) 1.56V d) 1.50V e) 1.47V
Answers:
1)
2)
3)
4)
5)
6)
7)
8)
9)
10)
11)
12)
d
e
a
b
a
e,
c
a
e
d
b
1997 B
In an electrolytic cell, a current of 0.250 ampere is passed through a solution of a chloride of iron,
producing Fe(s) and Cl2(g).
(a)
(b)
Write the equation for the half-reaction that occurs at the anode.
When the cell operates for 2.00 hours, 0.521 gram of iron is deposited at one electrode. Determine the
formula of the chloride of iron in the original solution.
(c) Write the balanced equation for the overall reaction that occurs in the cell.
(d) How many liters of Cl2(g), measured at 25˚C and 750 mm Hg, are produced when the cell operates as
described in part (b) ?
(e) Calculate the current that would produce chlorine gas from the solution at a rate of 3.00 grams per
hour.
Answer:
(a) 2 Cl– – 2 e- ∅ Cl2
(b)
0.250 amp ∞ 7200 sec = 1800 coulomb
1 mol e = 0.0187 mol e96500 coul
1 mol Fe
0.521 g Fe ∞
= 0.00933 mol Fe
55.85 g Fe
1800 coul ∞
0.0187 mol e 2
=
0.00933 mol Fe 1
Fe2+ + 2e- ∅ Fe ; ∴ FeCl2
(c)
Fe2+ + 2 Cl– ∅ Fe + Cl2
(d)
0.0187 mol e- ∞
1 mol Cl 2
= 0.00933 mol Cl2
2 mol e L • atm #
!
(0.00933 mol)" 0.0821
(298K)
nRT
mol • K $
V=
=
750
P
760 atm
(
)
= 0.231 L
(e)
3.00 g Cl 2
1 hr
1 mol Cl 2
×
×
1 hr
3600 sec 70.906 g Cl 2
∞
OR
2 mol e - 96500 amp⋅ sec
×
= 2.27 amp
1 mol Cl 2
1 mol e -
0.00933 mol Cl 2 0.662 g Cl 2 0.331 g Cl 2
=
=
2 hrs.
2 hrs.
1 hr.
0.250 amp
X
=
; X = 2.27 amp
0.331 g Cl 2 3.00 g Cl 2
Answer the following questions regarding the electrochemical cell shown.
(a) Write the balanced net-ionic equation for the spontaneous reaction that occurs as the cell operates, and
determine the cell voltage.
(b) In which direction do anions flow in the salt bridge as the cell operates? Justify your answer.
(c) If 10.0 mL of 3.0-molar AgNO3 solution is added to the half-cell on the right, what will happen to the
cell voltage? Explain.
(d) If 1.0 gram of solid NaCl is added to each half-cell, what will happen to the cell voltage? Explain.
(e) If 20.0 mL of distilled water is added to both half-cells, the cell voltage decreases. Explain.
Answer
(a) 2 Ag+ + 2 e- ∅ 2 Ag
E˚ = +0.80 v
Cd – 2 e-- ∅ Cd2+
2 Ag+ + Cd ∅ 2 Ag + Cd2+
(b)
(c)
(d)
E˚ = +0.40 v
E = +1.20v
Anions flow into the cadmium half-cell. As the cell operates, Cd2+ cations increase in number and
need to be balanced by an equal number of anion charges from the salt bridge.
Cell voltage will increase. An increase in silver ion concentration will result in faster forward reaction
and a higher cell potential.
Cell voltage will decrease. As the salt dissolves, the Cl– ion will cause the Ag+ ion to precipitate as
AgCl and decrease the [Ag+]. This will result in a slower forward reaction and a decrease in cell potential. Since cadmium chloride is a soluble salt, it will not affect the cadmium half-cell.
0.0592
[Cd 2 + ]
; while both concentrations are 1.0M, the cell potential is
log
2
[Ag + ]2
0.0592
[.5]
log
1.20v. But if each solution’s concentration is cut in half, then, E = 1.20v –
= 1.19v
2
[.5]2
(e)
Ecell = 1.20v –
cell
1993 D
A galvanic cell is constructed using a chromium electrode in a 1.00-molar solution of
Cr(NO3)3 and a copper electrode in a 1.00-molar solution of Cu(NO3)2. Both solutions
are at 25˚C.
(a) Write a balanced net ionic equation for the spontaneous reaction that occurs as the cell
operates. Identify the oxidizing agent and the reducing agent.
(b) A partial diagram of the cell is shown below.
Cr
Cu
1.0 M Cu(NO 3) 3
1.0 M Cr(NO 3)3
(i) Which metal is the cathode?
(ii) What additional component is necessary to make the cell operate?
(iii) What function does the component in (ii) serve?
(c) How does the potential of this cell change if the concentration of Cr(NO3)3 is changed
to 3.00-molar at 25˚C? Explain.
Answer:
(a) 2 Cr + 3 Cu2+ −−> 2 Cr3+ + 3 Cu
Cr = reducing agent; Cu2+ = oxidizing agent
(b) (i) Cu is cathode
(ii) salt bridge
(iii) transfer of ions or charge but not electrons
(c) Ecell decreases.
use the Nernst equation to explain answer