Unit 5 Electrochemistry S 2
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Transcript of Unit 5 Electrochemistry S 2
VERSION: April 20, 1999
Electrochemistry
A. Introduction
1. Electrochemistry is _______________________________________________________
________________________________________________________________________
________________________________________________________________________
Electrochemical reactions involve the _________________________________________
________________________________________________________________________
(analogous to proton transfer in acids and bases).
2. Consider the reaction
2Al(s) + 3CuCl2(aq) → AlCl3(aq) + 3Cu(s)
the net ionic equation is
This reaction only involves the ______________________________________________
We can re-write this equation as two separate _________________________
OXIDATION HALF-REACTION
REDUCTION HALF-REACTION
__________ of electrons is ____________________
__________ of electrons is ____________________
2 Chemistry 12
When a substance becomes oxidized _________________________________________
________________________________________________________________________
________________________________________________________________________
Zn → Zn2+ + 2e-
U3+ → U4+ + e-
2Cl- → Cl2 + 2e-
When a substance becomes reduced __________________________________________
________________________________________________________________________
________________________________________________________________________
Cu2+ + 2e- → Cu
V3+ + e- → V2+
F2 + 2e- → 2F-
3. For every reduction reaction there ____________________________________________
________________________________________________________________________
________________________________________________________________________
Reactions involving the loss and gain of electrons are called REDUCTION-
OXIDATION reactions or _______________ reactions.
In the reaction, 2Al + 3Cu2+ → 2Al3+ + 3Cu, Al is referred to as the _______________
________________________________________________________________________
________________________________________________________________________
while Cu2+ is referred to as the ______________________________________________
________________________________________________________________________
________________________________________________________________________
Unit 5 Electrochemistry (S).doc 3
The ____________________ _______________ is _______________ during a reaction.
The ____________________ _______________ is _______________ during a reaction.
B. Oxidation Numbers
1. An oxidation number is ____________________________________________________
________________________________________________________________________
________________________________________________________________________
2. Rules for determining oxidation number:
a) __________________________________________________________________
b) __________________________________________________________________
__________________________________________________________________
i) Li+, Na+, K+, and all other group 1 ions have _______________________
____________________________________________________________
ii) Ca2+, Ba2+, Mg2+, and all other group 2 ions have ___________________
____________________________________________________________
iii) F-, Cl-, Br-, I- (halogens) are usually ______________________________
___________________________________________________________ .
c) The oxidation number of _____________________________________________
In peroxides, H2O2, it is 1-.
d) The oxidation number for ____________________________________________
except in metallic hydrides such as NaH or BaH2 where it is 1-.
e) All other oxidation numbers are _______________________________________
__________________________________________________________________
__________________________________________________________________
4 Chemistry 12
EXAMPLE 5.1 DETERMINING OXIDATION NUMBERS
Problem: Determine the oxidation numbers of each atom for the following:
Na, CCl4, N2O4, PO43-, PbSO4
Solution: Na = _____
CCl4 = 0 = C + 4(_____)
⇒ Cl = _____, C = _____
N2O4 = 0 = 2N + 4(_____)
⇒ O = _____, N = _____
PO43- = 3- = P + 4(_____)
⇒ O = _____, P = _____
PbSO4 = 0 = Pb2+ + SO42-
⇒ Pb = _____
SO42- = 2- = S + 4(_____)
⇒ O = _____, S = _____
3. An atom’s change in oxidation number indicates whether it is oxidized or reduced. When
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
ClO3- → ClO4
- (_____ → _____ ; _______________)
H2O2 → H2O (_____ → _____ ; _______________)
Cr3+ → CrO42- (_____ → _____ ; _______________)
NO2 → N2O3 (_____ → _____ ; _______________)
Unit 5 Electrochemistry (S).doc 5
OXIDATION = __________________________________________________________
REDUCTION = _________________________________________________________
C. Predicting Spontaneous Reactions
1. An excerpt from the Table of Reduction Potentials is shown below.
F2 + 2e- →← 2F -
Ag+ + e- →← Ag
Cu2+ +2e- →← Cu
Zn2+ + 2e- →← Zn
Li + + e- →← Li
INCREASING TENDENCY TO REDUCE
INCREASING STRENGTH
AS AN OXIDIZING AGENT
INCREASING TENDENCY TO OXIDIZE
INCREASING STRENGTH
AS A REDUCING AGENT
The following observation can be made from the Table of Reduction Potentials:
i) In general, __________________________________________________
____________________________________________________________
____________________________________________________________
ii) In general, __________________________________________________
____________________________________________________________
____________________________________________________________
iii) Some metals such as Fe, Sn, Cr, Hg, and Cu ________________________
____________________________________________________________
As a result, __________________________________________________
____________________________________________________________
____________________________________________________________
Remember this point because it means that _________________________
6 Chemistry 12
____________________________________________________________
e.g. Cu+ + e- →← Cu(s)
Cu2+ + 2e- →← Cu(s)
Cu2+ +e- →← Cu+
The Cu+ ion is found on both left and right sides. Similarly, Sn2+ and
Fe2+ are also found on both sides of the table.
iv) ____________________________________________________________
____________________________________________________________
____________________________________________________________
OXIDIZING AGENTS ___________________________________________________
________________________________________________________________________
REDUCING AGENTS ___________________________________________________
________________________________________________________________________
2. The table of Standard Reduction Potentials is used in a similar way to the table of
Relative Strengths of Acids. The species at the _______________ __________ have a
tendency to go _______________ while the species at the _______________
__________ have a tendency to go _______________. Each reaction in the table can go
________________________________________________________________________
The half-reaction for Zn and Zn2+ is
If there is a piece of Zn(s) in a solution of Zn2+, you can write either:
Zn2+ + 2e- → Zn (written as a _______________)
or Zn → Zn2+ + 2e- (written as an _______________)
Unit 5 Electrochemistry (S).doc 7
Note: When referring to an isolated half-reaction, use equilibrium arrows to show that
the reaction can go forward or backward.
Ag+ + e- →← Ag
If the half-reaction is made to undergo either reduction or oxidation as a result of
being part of a redox reaction, the use a one-way arrow.
Ag+ + e- → Ag
3. In a redox reaction, _______________________________________________________
________________________________________________________________________
________________________________________________________________________
Assume that you have Zn in a solution of Zn2+ and Cu in a solution of Cu2+, the two
possible half reactions are
Cu2+ + 2e- →← Cu
and Zn2+ + 2e- →← Zn
Of the two oxidizing agents, _____ and _____, the _____ is _______________ on the
left side of the table and has a greater tendency to become _______________. The
reduction reaction will be
Cu2+ + 2e- → Cu
8 Chemistry 12
Of the two reducing agents, _____ and _____, the _____ is _______________ on the
right side of the table and has a greater tendency to become _______________. The
oxidation reaction will be
Zn2+ + 2e- ← Zn
or re-written as
Zn → Zn2+ + 2e-
The overall redox reaction which will spontaneously occur is found by adding the two
half-reactions. Cu2+ + 2e- → Cu
Zn → Zn2+ + 2e-
When two half-reactions are joined, ________________________________________
________________________________________________________________________
________________________________________________________________________
4. If you are _____________________________________________, rather than the four
necessary for two complete half-reactions as above, you __________________________
________________________________________________________________________
a) If ________________________________________________________________
__________________________________________________________________
(i.e., both are oxidizing agents (left) or both reducing agents (right)) then
__________________________________________________________________
Unit 5 Electrochemistry (S).doc 9
Assume the only reactants are Zn and Cu
Cu2+ + 2e- ← Cu
Zn2+ + 2e- ← Zn
Both Cu and Zn can only undergo
_______________, there is nothing present to
undergo _______________. A complete redox
reaction is not possible. ___________________
Assume the only reactants are Zn2+ and Cu2+
Cu2+ + 2e- → Cu
Zn2+ + 2e- → Zn
Both Cu2+ and Zn2+ can only undergo
_______________, there is nothing present to
undergo _______________. A complete redox
reaction is not possible. ___________________
b) If __________________________________________________ (oxidizing
agent) and __________________________________________________
(reducing agent) there are two possible cases.
i) The ____________________ _______________ is _______________ on
the table than the reducing agent.
Assume the reactants are Cu2+ and Zn. Cu2+ + 2e- → Cu
Zn2+ + 2e- ← Zn
The overall redox reaction is obtained by reversing the oxidation reaction
and adding the two half-reactions. Cu2+ + 2e- → Cu
Zn → Zn2+ + 2e-
This is a ____________________ _______________
10 Chemistry 12
ii) The ____________________ _______________ is _______________ on
the table than the reducing agent.
Assume the reactants are Zn2+ and Cu. Cu2+ + 2e- → Cu
Zn2+ + 2e- ← Zn
In this case, _____ ____________________ occurs.
Notice that this is the reverse reaction to (i). Since the reaction
Cu2+ + Zn → Cu + Zn2+
is ____________________ , the reverse reaction is
_________________________
A reaction will be _________________________________________________________
________________________________________________________________________
(reactant to be reduced – left) which is _________________________________________
______________________________________________ (reactant to be oxidized – right).
Unit 5 Electrochemistry (S).doc 11
EXAMPLE 5.2 PREDICTING SPONTANEOUS REACTIONS
Problem: Predict whether or not a reaction will occur when the following are mixed.
a) Cl2 and Br-
b) Sn and Mn
c) Ni2+ and Pb
Solution: a) Cl2 and Br-
Reaction is _________________________and the overall reaction is
b) Sn and Mn
Both reactants are ____________________________________________
___________________________________________________________
c) Ni2+ and Pb
___________________________________________________________
___________________________________________________________
12 Chemistry 12
5. In _____________________________________________________________________
________________________________________________________________________
When H+ is present, _______________________________________________________
For example, if you asked whether the reaction of SO42- in a solution of Na2SO4 will
reduce:
SO42- + 4H+ + 2e- → H2SO3 + 2H2O …… +0.17
your answer should be _____________________________________________________
________________________________________________________________________
________________________________________________________________________
Note: H+ is necessary in many reduction half-reactions; however, there is only ONE
reaction is which H+ is the only substance involved in the reduction half-reaction:
2H+ + 2e- → H2(g) …… 0.00
Unit 5 Electrochemistry (S).doc 13
D. Balancing Half Reactions
1. A half-reaction ___________________________________________________________
When balancing half-reactions ______________________________________________
________________________________________________________________________
The steps to balancing a half-reaction are as follows:
Typically you will be starting with a “skeleton equation” containing the major atoms
involved.
a) Balance ___________________________________________________________
__________________________________________________________________
b) Balance ___________________________________________________________
__________________________________________________________________
c) Assume that solutions are acidic and ____________________________________
__________________________________________________________________
d) Balance ___________________________________________________________
__________________________________________________________________
14 Chemistry 12
EXAMPLE 5.3 BALANCING HALF-REACTIONS IN ACIDIC SOLUTION
Problem: Balance the following half-reaction:
Cr2O72- → Cr3+
Solution: Step 1: Balance ATOMS.
Cr2O72- → ____Cr3+
Step 2: Balance OXYGEN.
Cr2O72- → 2Cr3+ + ________
Step 3: Balance HYDROGEN.
________ + Cr2O72- → 2Cr3+ + 7H2O
Step 4: Balance CHARGE.
________ + 14H+ + Cr2O72- → 2Cr3+ + 7H2O
Note: Electrons are negatively charged, ___________________________
______________________________________________________
2. Sometimes it will be necessary to balance half-reactions in BASIC solutions. For these
problems, _______________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
Unit 5 Electrochemistry (S).doc 15
EXAMPLE 5.3 BALANCING HALF-REACTIONS IN BASIC SOLUTION
Problem: Balance the following half-reaction:
MnO4- → MnO2
Solution: Step 1: Balance ATOMS.
MnO4- → MnO2
Atoms already balanced.
Step 2: Balance OXYGEN.
MnO4- → MnO2 + ________
Step 3: Balance HYDROGEN.
________ + MnO4- → MnO2 + 2H2O
Step 4: Balance CHARGE.
________+ 4H+ + MnO4- → MnO2 + 2H2O
Step 5: Add OH- and cancel out H2O.
________+ 3e- + 4H+ + MnO4- → MnO2 + 2H2O + ________
3e- + ________ + MnO4- → MnO2 + 2H2O + ________
3e- + ________ + MnO4- → MnO2 + ________
16 Chemistry 12
E. Balancing Redox Equations Using Half-Reactions
1. An overall redox equation can be obtained by __________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
EXAMPLE 5.4 BALANCING A REDOX REACTION BY THE HALF-REACTION METHOD
Problem: Balance the following half-reaction:
ClO4- + I2 → Cl
- + IO3
-
Solution: Step 1: Break equation into two half-reactions.
________________________
________________________
Step 2: Balance each half-reaction.
________________________________________
________________________________________
Step 3: Lowest common multiple between electrons.
In this case LCM is ________, multiply reduction by ____ and oxidation by ____.
5 x (8e- + 8H+ + ClO4- → Cl
- + 4H2O)
4 x (6H2O + I2 → 2IO3- + 12H+ + 10e-)
____e- + ____H+ + ____ClO4- → ____Cl
- + ____H2O
____H2O + ____I2 → ____IO3- + ____H+ + ____e-
Unit 5 Electrochemistry (S).doc 17
Step 4: Add two half-reactions, cancelling out electrons and common particles.
40e- + 40H+ + 5ClO4- → 5Cl
- + 20H2O
24H2O + 4I2 → 8IO3- + 48H+ + 40e-
______________________________________
Balanced redox equations ___________________________________________________
________________________________________________________________________
________________________________________________________________________
2. Once again, in basic solutions the final equation can be converted by ________________
________________________________________________________________________
________________________________________________________________________
5ClO4- + 4I2 + 4H2O → 5Cl- + 8IO3
- + 8H+
For basic solution, add 8OH- to both sides of equation
____ + 5ClO4- + 4I2 + 4H2O → 5Cl- + 8IO3
- + 8H+ + ____
____ + 5ClO4- + 4I2 + 4H2O → 5Cl- + 8IO3
- + ____
Cancel out water
8OH- + 5ClO4- + 4I2 + 4H2O → 5Cl- + 8IO3
- + 8H2O
8OH- + 5ClO4- + 4I2 → 5Cl- + 8IO3
- + ________
18 Chemistry 12
3. In some redox reactions, ___________________________________________________
________________________________________________________________________
________________________________________________________________________
Such a reaction is called a ________________________
ClO2- → ClO3
- + Cl-
The two half-reactions are
ClO2- → ClO3
-
ClO2- → Cl-
Balanced redox equation is
3ClO2- → 2ClO3
- + Cl-
F. Balancing Redox Equations Using Oxidation Numbers
1. This method involves ______________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
Remember that an increase in oxidation number is oxidation and a decrease in oxidation
number is reduction. In this method __________________________________________
________________________________________________________________________
________________________________________________________________________
Unit 5 Electrochemistry (S).doc 19
EXAMPLE 5.5 BALANCING A REDOX REACTION BY OXIDATION NUMBERS
Problem: Balance the following half-reaction:
ClO4- + I2 → Cl
- + IO3
-
Solution: Step 1: Balance atoms.
ClO4- + I2 → Cl
- + 2IO3
-
Step 2: Assign oxidation numbers to all atoms involved in a change in oxidation number (∆ON = “change in oxidation number”).
ClO4- + I 2 ⎯⎯→ Cl
- + 2IO 3
-
Step 3: Balance change in oxidation number
In this case LCM is ____, multiply decrease by ____ and increase by ____.
____ClO4- + ____I2 ⎯⎯→ ____Cl
- + ____IO3
-
Step 4: Balance OXYGENS by adding H2O and HYDROGENS by adding H+. DO NOT ADD ELECTRONS.
________ + 5ClO4- + 4I2 ⎯⎯→ 5Cl
- + 8IO3
- + ________
Remember, ______________________________________________________________
________________________________________________________________________
________________________________________________________________________
20 Chemistry 12
SAMPLE
PROBLEM 5.1 BALANCING A REDOX REACTION BY OXIDATION NUMBERS
Problem: Balance the following redox reactions by the oxidation number method:
a) MnO4- + Fe2+ → Mn2+ + Fe3+ (acidic)
b) S2- + ClO3- → Cl- + S (basic)
c) CN- + IO3- → I- + CNO- (basic)
d) H2O2 + Cr3+ → CrO42- + H2O (acidic)
Solution:
MnO4- + Fe2+ → Mn2+ + Fe3+
S2- + ClO3- → Cl - + S
CN - + IO3- → I - + CNO -
H2O2 + Cr3+ → CrO42- + H2O
Unit 5 Electrochemistry (S).doc 21
G. Redox Titrations
1. Acid-base titrations are useful for determining the concentration of an unknown sample
of acid or base. In a similar manner, __________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
2. OXIDIZING AGENTS
One of the most common oxidizing agents used in redox titrations is ________________
The half-reaction:
MnO4- + 8H+ + 5e- → Mn2+ + 4H2O ……… E° = 1.49 V
has such a strong tendency to reduce that it will be able to oxidize a large number of
other substances. (The K+ in KMnO4 is left out because it is a spectator ion.)
In acid-base titrations, an ____________________ which changes colour is added to help
identify the ____________________ _______________ of the titration. An indicator is
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
If MnO4- (oxidizing agent) in a burette is added to a solution containing Fe2+ (reducing
agent), the MnO4- is converted to Mn2+ (colourless). ____________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
22 Chemistry 12
EXAMPLE 5.6 REDOX TITRATION INVOLVING THE PERMANGANATE ION
Problem: A 100.0 mL sample containing FeCl2 is titrated with 0.100 M KMnO4 solution. If 27.45 mL of KMnO4 solution were required to reach the end-point, what was the [FeCl2]?
Solution: First we need the balanced redox equation.
Fe2+
+ MnO4- → Fe
3+ + Mn
2+
___ Fe2+
+ MnO4- + _______ → ___Fe
3+ + Mn
2+ + _______
The remainder of the calculation is similar to acid-base titration calculations.
Calculate MOLES of MnO4-
0.100 M MnO4- x
27.45 mL1000 mL = _______________ mol MnO4
-
Use MOLE RATIO from balance equation.
0.002745 mol MnO4- x
5 mol Fe2+
1 mol MnO4- = _______________ mol Fe
2+
Determine CONCENTRATION.
0.0137 mol Fe2+
0.100 L = _______________M Fe2+
and since Fe2+ and FeCl2 are in a 1:1 ratio, [FeCl2] = _______________.
Unit 5 Electrochemistry (S).doc 23
3. REDUCING AGENTS
A commonly used reducing agent is __________ or __________. A large number of
substances can oxidize I- to I2 according to the half-reaction
2I- → I2 + 2e-
Titrations involving I- generally involve two consecutive steps:
a) first, _____________________________________________________________
__________________________________________________________________
b) second, ___________________________________________________________
__________________________________________________________________
__________________________________________________________________
An example of a reaction involving I- is the reduction of laundry bleach, NaOCl. The
reaction between I- and OCl- proceeds according to:
2I- → I2 + 2e-
2e- + 2H+ + OCl- → Cl- + H2O ___________________________________
No attempt is made to add exactly enough I- to react with the OCl-. Rather, __________
________________________________________________________________________
The excess I- does not affect the results; however, _______________________________
________________________________________________________________________
The above reaction between OCl- and I- is the “initial” reaction because the actual redox
titration involves a second reaction between the I2 produced and the reducing agent
sodium thiosulphate, Na2S2O3.
I2 + 2e- → 2I-
2S2O32- → S4O6
2- + 2e- ___________________________
24 Chemistry 12
When the addition of S2O32- has reacted most of the I2 present, the brown colour of I2
almost disappears (pale yellow). Some starch solution is then added to the titration,
producing a dark blue colour (from the reaction of starch with the remaining I2 in
solution). After adding the starch, ___________________________________________
________________________________________________________________________
________________________________________________________________________
Unit 5 Electrochemistry (S).doc 25
EXAMPLE 5.7 REDOX TITRATION INVOLVING THE IODIDE ION
Problem: A 25.00 mL sample of bleach is reacted with excess KI according to the equation:
2H+ + OCl- + 2I- → Cl- + H2O + I2
The I2 produced requires exactly 46.84 mL of 0.7500 M Na2S2O3 to bring the titration to the endpoint according to the equation:
2S2O32- + I2 → S4O6
2- + 2I-
using starch solution as an indicator. What is the [OCl-] in the bleach?
Note: The TWO reactions are linked together. The I2 produced in the first reaction is used up in the second
2H+ + OCl- + 2I- → Cl- + H2O + I2
2S2O32- + I2 → S4O6
2- + 2I-
Solution: First, calculate the moles of S2O32-
moles S2O32- = __________ M x
46.84 mL1000 mL/L = _______________ mol S2O3
2-
Next, calculate the moles of I2
moles I2 = __________ mol S2O32- x
I22S2O3
2- = _______________ mol I2
Now moles of OCl - can be calculated
moles OCl - = __________ mol I2 x OCl -
I2 = __________ mol OCl -
Finally, the [OCl -] can be calculated
[OCl -] = 0.01757 mol0.02500 L = _______________
26 Chemistry 12
H. Electrochemical Cells
1. An _________________________ __________ (galvanic or voltaic cell) is a chemical
system in which __________________________________________________________
________________________________________________________________________
Consider the reaction
Cu2+ + Zn → Cu + Zn2+
A spontaneous reaction will occur when zinc metal is placed into a solution of CuSO4;
however, ________________________________________________________________
________________________________________________________________________
This same reaction can used to produce electricity if _____________________________
________________________________________________________________________
2. An electrochemical cell:
a) In the electrochemical cell above, we have Zn(s), Zn(NO3)2(aq), Cu(s), and
CuSO4(aq). The half-reactions are:
Cu2+ + 2e- →← Cu …… + 0.34
Zn2+ + 2e- →← Zn …… - 0.76
Unit 5 Electrochemistry (S).doc 27
Since Cu2+ is a ______________________________________ than Zn2+, Cu2+
will become ____________________ and since Zn is a _____________________
than Cu, Zn will become ____________________. The reduction and oxidation
half-reactions that occur will be →
→ __________________________
Cu2+ + Zn → Cu + Zn2+
28 Chemistry 12
b) In the copper cell, ___________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
At the copper electrode, copper ions, Cu2+, ______________________________
__________________________________________________________________
The copper electrode is ______________________________________________
__________________________________________________________________
__________________________________________________________________
Cu2+ + 2e- → Cu
In the zinc cell, _____________________________________________________
___________________________________. At the zinc electrode, zinc metal ___
The zinc electrode is ________________________________________________
__________________________________________________________________
Zn → Zn2+ + 2e-
Unit 5 Electrochemistry (S).doc 29
_______________ is the electrode where ____________________ occurs.
_______________ is the electrode where ____________________ occurs.
c) The two electrodes are connected by a __________________________________
_________________________. Electrons produced at the zinc electrode flow
through the wire to the copper electrode where they reduce copper ions. The
electrons will ______________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
A __________ _______________ composed of a concentrated solution of a
strong electrolyte, usually KNO3 or KCl, connects the two cells to complete the
circuit. As the two half-reactions occur, _________________________________
__________________________________________________________________
__________________________________________________________________
Water and certain ions, Zn2+ and SO42-, will pass through the salt bridge from
one cell to the other. The positive Zn2+ ions are attracted to the CATHODE so
they are called _______________ while the negative SO42- ions are attracted to
the ANODE so they are called _______________.
__________________________________________________________________
__________________________________________________________________
30 Chemistry 12
Note: There are NO electrons flowing in the solution, only ions. ____________
____________________________________________________________
The salt bridge allow __________________________________________
____________________________________________________________
Since oxidation occurs at the anode, ______________________________
and since reduction occurs at the cathode, the _______________________
__________ as the electrochemical cell operates.
Unit 5 Electrochemistry (S).doc 31
EXAMPLE 5.8 DRAWING AN ELECTROCHEMICAL CELL
Problem: Assume two half-cells consisting of Pb(s) in a Pb(NO3)2 solution and Zn(s) in a ZnCl2 solution are connected to make an electrochemical cell. Draw and label the parts of the cell, write the equations for the individual half-reactions and overall reaction, and indicate the directions in which the ions and electrons move.
Solution: The two half-reactions are:
Pb2+ + 2e- →← Pb …… - 0.13
Zn2+ + 2e- →← Zn …… - 0.76
Since Pb2+ is the stronger oxidizing reagent, it will become reduced and Zn(s) will become oxidized.
The two half-reactions are
Pb2+ + 2e- → Pb
Zn → Zn2+ + 2e- __________________________
Pb2+
+ Zn → Pb + Zn2+
32 Chemistry 12
I. Standard Reduction Potentials
1. The tendency of electrons to flow in an electrochemical cell is called the
____________________, or ____________________ ____________________to do
work.
________________________________________________________________________
________________________________________________________________________
Since electrons cannot flow in an isolated half-cell, ______________________________
________________________________________________________________________
However, the ____________________________________________________________
________________________________________________________________________
A ______________________________________________________________________
________________________________________________________________________
The voltage for the HYDROGEN HALF-CELL is defined to be
2H+ + 2e- → H2 …… E° = __________
Where E° is the _____________________________________________, in volts and the
“ ° ” in E° means _________________________ ____________________
An electrochemical cell is said to be at STANDARD STATE if
• it is at __________, and
• all gases are at _________________________, and
• all elements are in their standard states (normal phases at 25 °C), and
• all solutions involved in the cell have a ______________________________
Unit 5 Electrochemistry (S).doc 33
2. All voltages listed in the table of STANDARD REDUCTION POTENTIALS are
determined at standard state and are _________________________________________
________________________________________________________________________
Cu2+ + 2e- →← Cu …… E° = + 0.34
This half-cell has a voltage which is 0.34 V ____________________ than that of the
hydrogen half-cell.
Zn2+ + 2e- →← Zn …… E° = - 0.76
This half-cell has a voltage which is 0.76 V __________ than that of the hydrogen half-
cell.
Since the voltage is a measure of the work done, reversing a reduction reaction such as
Zn2+ + 2e- → Zn …… E° = - 0.76
produces ________________________________________________________________
Zn → Zn2+ + 2e- …… E° = + 0.76
________________________________________________________________________
________________________________________________________________________
3. When two half-reactions are combined, _______________________________________
________________________________________________________________________
________________________________________________________________________
Hg2+ + 2e- →← Hg …… E° = + 0.85 V
Cu2+ + 2e- →← Cu …… E° = + 0.34 V
34 Chemistry 12
The voltage for the overall reaction is
→ …… E° = __________ V
→ …… E° = __________ V _______________________________________________
→ …… E°CELL = __________ V
The potential of an electrochemical cell is ____________________________________
________________________________________________________________________
________________________________________________________________________
EXAMPLE 5.9 CALCULATING CELL POTENTIALS
Problem: Calculate the potential of the cell: Ni2+ + Fe → Ni + Fe2+.
Solution: The half-reactions involved are:
Ni 2+ + 2e- →← Ni …… E° = - 0.26 V
Fe2+ + 2e- →← Fe …… E° = - 0.45 V
The voltage for the overall reaction is
→ …… E° = __________V
→ …… E° = __________ V _________________________________________ → …… E°CELL = __________
Note: This reaction is SPONTANEOUS since Ni2+ is higher than Fe and the reaction has a POSITIVE VOLTAGE.
Unit 5 Electrochemistry (S).doc 35
EXAMPLE 5.10 CALCULATING CELL POTENTIALS
Problem: Calculate the potential of the cell: Ni + Fe2+ → Ni 2++ Fe.
Solution: The half-reactions involved are:
Ni 2+ + 2e- →← Ni …… E° = - 0.26 V
Fe2+ + 2e- →← Fe …… E° = - 0.45 V
The voltage for the overall reaction is
→ …… E° = __________ V
Fe2+ + 2e- → Fe …… E° = __________ V _______________________________________________
→ …… E°CELL = __________
Note: This reaction is NOT SPONTANEOUS since Fe2+ is lower than Ni and the reaction has a NEGATIVE VOLTAGE.
EXAMPLE 5.11 CALCULATING CELL POTENTIALS
Problem: Calculate the potential of the cell: 3Ag+ + Al → 3Ag + Al3+.
Solution: The half-reactions involved are:
Ag+ + e- →← Ag …… E° = + 0.80 V
Al3+ + 3e- →← Al …… E° = - 1.66 V
The voltage for the overall reaction is
→ …… E° = __________ V
→ …… E° = __________ V _______________________________________________
→ …… E°CELL = __________
Note: Although the Ag+/Ag half-reaction is multiplied by 3 to balance the electrons, the voltage is NOT multiplied by 3.
36 Chemistry 12
If E° is _______________ for a redox reaction, the reaction is expected to be
____________________
If E° is ____________________ for a redox reaction, the reaction is
_________________________
Note: Although E° can be used to predict if a reaction is spontaneous, it ___________
__________________________________________________________________
__________________________________________________________________
Remember the ____________________ _______________ of a reaction
determines the rate.
In some texts, the reduction of neutral water is given as
2H+ (10-7 M) + 2e- →← H2(g) …… E = - 0.41 V
Other texts show this same reaction as
2H2O + 2e- →← H2(g) + 2OH- (10-7 M) …… E = - 0.41 V
If a redox reaction occurs in acidic water, the reduction of H+ (at 0.00 V) may
occur and if a redox reaction occurs in a neutral solution, the reduction of neutral
water (at - 0.41 V) may be possible.
4. The SURFACE AREA of an electrode _______________________________________
________________________________________________________________________
This is because ___________________________________________________________
________________________________________________________________________
Therefore, increasing the surface area of a solid electrode has no effect on the
concentration of the solid. As a result there is no equilibrium shift and no change in half-
cell potential.
Unit 5 Electrochemistry (S).doc 37
Increasing the surface are of an electrode does however, __________________________
____________________. The number of electrons transferred per second increases.
This increases the ____________________ of the cell but not the voltage. Increasing the
surface area of the electrode, also _____________________________________________
________________________________________________________________________
5. When a half-reaction is NOT at standard state, it is ______________________________
________________________________________________________________________
Consider the following half-reaction
Cu2+ + 2e- →← Cu …… E° = + 0.34 V
If the [Cu2+] is _______________ than 1.0 M the equilibrium shifts __________ and the
reduction potential ____________________.
Cu2+ + 2e- →← Cu …… E > + 0.34 V
If the [Cu2+] is __________ than 1.0 M the equilibrium shifts __________ and the
reduction potential ____________________.
Cu2+ + 2e- →← Cu …… E < + 0.34 V
Notice that the “ ° ” is omitted from the E° symbol since it is not at standard state.
6. Operating electrochemical cells ______________________________________________
Consider the cell
2Ag+ + Cu → 2Ag+ Cu2+
THE REDUCTION REACTION: 2Ag+ + 2e- → 2Ag
As the [Ag+] decreases, ____________________________________________________
________________________________________________________________________
As a result, the ___________________________________________________________
________________________________________________________________________
38 Chemistry 12
THE OXIDATION REACTION: Cu2+ + 2e- ← Cu
As the [Cu2+] increases, ____________________________________________________
________________________________________________________________________
(the reduction potential of Cu2+ + 2e- → Cu increases) Therefore, the ______________
________________________________________________________________________
________________________________________________________________________
Overall, the following occurs as the cell goes to equilibrium.
2Ag+ + 2e- → 2Ag
Cu2+ + 2e- ← Cu
Voltage 2Ag+ + 2e- → 2Ag
Cu2+ + 2e- ← Cu
INITIALLY AS TIME PASSES
FINALLY
E° = + 0.46 V
E = 0.00 V
CELL
CELL
________________________________________________________________________
________________________________________________________________________
Unit 5 Electrochemistry (S).doc 39
J. Selecting Preferred Reactions
1. When a cell contains a mixture of substances, __________________________________
________________________________________________________________________
Consider the following cell:
Zn AgCu
Zn
NO
AgCu
3
2++
2+-
NO3-
The possible half-reactions are:
_____________________________________________
_____________________________________________
_____________________________________________
When several different reduction half-reactions can occur, _____________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
Therefore, the __________ will reduce and form __________
40 Chemistry 12
When several different oxidation half-reactions can occur, _____________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
Therefore, the __________ will oxidize to form __________
2. Determining the preferred half-reactions:
• First, _____________________________________________________________
__________________________________________________________________
• Next, starting at the _________________________________________________
__________________________________________________________________
__________________________________________________________________
• Then, starting at the _________________________________________________
__________________________________________________________________
__________________________________________________________________
Unit 5 Electrochemistry (S).doc 41
EXAMPLE 5.12 SELECTING PREFERRED REACTIONS
Problem: An iron strip is placed in a mixture of Br2(aq) and I2(aq). What is the preferred reaction which occurs?
Solution: The species present and the possible reactions are
Br2 + 2e- →← 2Br- …… E° = + 1.09 V
I2 + 2e- →← 2I- …… E° = + 0.54 V
Fe2+ + 2e- →← Fe …… E° = - 0.45 V
The preferred reduction involves __________ and the only oxidation possible involves __________. The overall reaction is
→
→
___________________________
→
Note: Any ion capable of being reduced ______________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
Similarily, any ion capable of being ____________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
The following ions are generally considered to be spectator ions:
Na+, K+, Ca2+, Mg2+, SO42- (in neutral solution), and Cl-
42 Chemistry 12
K. Corrosion of Metals
1. The term CORROSION simply _____________________________________________
The most common form of corrosion _________________________________________
When a drop of water rests on an iron surface, a spontaneous reaction occurs.
Fe
H O
cathode region (oxygen-rich)
Precipitate of Fe(OH)
anode region (oxygen-poor)
2
2
At the __________________________________________________________________
________________________________________________________________________
Fe(s) → Fe2+ + 2e-
After being formed, the Fe2+ ions ____________________________________________
________________________________________________________________________
________________________________________________________________________
As the Fe2+ ions move away, they expose fresh Fe(s) for further oxidation. In the
meantime, the reaction 12 O2 + H2O + 2e- → 2OH-
is proceeding at the oxygen-rich outer surface of the drop. When the Fe2+ meets the OH-
, it _____________________________________________________________________
The Fe(OH)2 is eventually __________________________________________________
________________________________________________________________________
Unit 5 Electrochemistry (S).doc 43
“Rust” is just Fe2O3•XH2O, where “X” is variable. The variability of “X” explains the
different colours rust can have (red, brown, yellow, black) since differing numbers of
water molecules attached to the Fe2O3 will change the colour of the component.
A metal _________________________________________________________________
________________________________________________________________________
________________________________________________________________________
For example, if iron touches some copper wire and the place where they touch gets wet,
then:
Fe(s) → Fe2+ + 2e- (Fe has a greater tendency to oxidize than Cu)
The copper ______________________________________________________________
________________________________________________________________________
________________________________________________________________________ 12 O2 + H2O + 2e- → 2OH-
2. PREVENTING CORROSION
A. Isolating The Metal From The Environment
i) Apply ______________________________________________________
____________________________________________________________
If oxygen and water cannot reach the metal, it will not corrode.
ii) Apply ______________________________________________________
____________________________________________________________
Tin can be applied to the surface of steel cans. The tin is quickly oxidized
to produce a thin layer of tin oxide which protects the underlying metal
from further corrosion.
44 Chemistry 12
B. Electrochemical Methods
i) Cathodic Protection
Recall that __________________________________________________
____________________________________________________________
____________________________________________________________
____________________________________________________________
This implies that we can prevent the corrosion of iron ________________
____________________________________________________________
____________________________________________________________
The term CATHODIC PROTECTION is applied to the process of ____
____________________________________________________________
____________________________________________________________
____________________________________________________________
For example, both ____________________ and __________ are stronger
reducing agents than iron, if pieces of magnesium or zinc are attached to
the surface of iron, the magnesium and zinc will ____________________
____________________________________________________________
Since the substance with the greater tendency to oxidize, Mg, acts as the
_______________, the Fe acts as a _______________ and remains in its
reduced form.
e.g. Zinc blocks are bolted to iron-hulled ships to prevent them from rusting.
Some modern ships pass a low voltage electric current into the hull from
an electrical generator. This forces electrons into the metal and prevents it
from becoming oxidized.
Galvanized iron involves coating zinc onto the surface of iron sheets.
Unit 5 Electrochemistry (S).doc 45
Some buried gas and oil tanks made from steel have thick braided wire
connected to them. The wire is attached to posts made of magnesium or
zinc.
ii) Change the Condition of the Surroundings
When iron is placed in contact with water containing oxygen, the half-
reaction 12 O2 + H2O + 2e- → 2OH-
oxidizes the iron. If _______________ is removed from the solution, the
reduction tendency drops drastically.
In addition to removing the oxygen, ______________________________
____________________________________________________________
____________________________________________________________
L. Electrolysis
1. ELECTROLYSIS is ______________________________________________________
________________________________________________________________________
________________________________________________________________________
An ____________________ __________ or ELECTROLYSIS CELL is an apparatus
in which electrolysis can occur.
____________________ __________ involve ____________________ redox reactions
that require energy to occur ____________________
____________________ __________ involve ____________________ redox reactions
that release energy ____________________
46 Chemistry 12
2. ELECTROLYSIS OF A MOLTEN BINARY SALT
A _______________ __________ is made up of only two different elements, e.g. NaCl,
KBr, MgI2, AlF3, etc. When a salt such as NaCl is melted, the ions are free to move
(Note: molten NaCl is NaCl(l), there are only Na+ and Cl- ions present. Do not confuse
with NaCl(aq) which is a solution and contains Na+, Cl-, and H2O).
– +
Na Cl+
–
direct current (DC) power supply or battery.
(the symbol is sometimes used)
There is ________________________________________________________________
________________________________________________________________________
The only reactants present are Na+ and Cl-
Cl2 + 2e- → 2Cl-
Na+ + e- → Na
The _______________ reaction is 2Cl- → Cl2 + 2e- E° = - 1.36 V
The _______________ reaction is Na+ + e- → Na E° = - 2.71 V
The overall reaction is 2Na+ + 2Cl- → 2Na + Cl2 E°cell = - 4.07 V
Note: This is a _________________________ reaction since the oxidizing agent is
below the reducing agent and since the voltage is negative.
To make this cell operate, ___________________________________________________
Unit 5 Electrochemistry (S).doc 47
Since the half-cells are not at standard state, the reduction potentials will be different
from those listed on the table.
3. ELECTROLYSIS OF AN AQUEOUS SOLUTION
Consider the electrolysis of aqueous sodium iodide, NaI(aq). This __________________
________________________________________________________________________
Once again, _______________ electrodes are used and the cell is set up as follows,
– +
D.C.
Note: During the electrolysis of aqueous solutions, _____________________________
__________________________________________________________________
There are two possible oxidations: 12 O2(g) + 2H+ (10-7 M) + 2e- ←
I2 + 2e- ←
E� = + 0.82 V
E� = + 0.54 V
The oxidation of __________________________________________________________
________________________________________________________________________
48 Chemistry 12
There are two possible reductions:
+ 2e- → H2(g) + 2OH- (10-7 M)
+ e- → Na
E� = - 0.41 V
E� = - 2.71 V
The reduction of __________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
________________________________________________________________________
The half-reactions and the overall reaction involving the electrolysis of NaI(aq) are
2H2O + 2e- → H2(g) + 2OH- (10-7 M) E� = - 0.41 V
2I- → I2 + 2e- E� = - 0.54 V
E� = - 0.95 V
so at least + 0.95 V must be applied to the cell.
Unit 5 Electrochemistry (S).doc 49
Note: The concentrations of the materials in cells will not be of concern; as long as
there is sufficient material in the cell you can assume the reactions proceed as
predicted.
In NEUTRAL aqueous solutions, there are two “water equations” which must be
considered.
OXIDATION: 12 O2(g) + 2H+ (10-7 M) + 2e- ← H2O E� = + 0.82 V
REDUCTION: 2H2O + 2e- → H2(g) + 2OH- (10-7 M) E� = - 0.41 V
In ACIDIC solutions there are also two “water equations” which must be
considered.
OXIDATION: 12 O2(g) + 2H+ + 2e- ← H2O E� = + 1.23 V
REDUCTION: 2H+ + 2e- → H2(g) E� = 0.00 V
4. In practice, it is found that a higher potential than calculated must be applied to cause
electrolysis. The causes are _________________________________________________
________________________________________________________________________
The difference between actual potentials required for electrolysis and the calculated
potential is termed ___________________________________. Overpotentials vary for
each half-cell but the difference in potentials is _________________________________
________________________________________________________________________
As a result of the overpotential effect, when dilute neutral solutions (< 1.0 M) containing
_____ or _____ are electrolyzed
50 Chemistry 12
OXIDATION:
Cl2 + 2e- ← 2Cl - E� = + 1.36 V
12 O2(g) + 2H+ (10-7 M) + 2e- ← H2O E� = + 0.82 V
Br2 + 2e- ← 2Br - E� = + 1.09 VStrength of Reducing
Agent
we would expect that _____ will be produced. In practice however, we find that _____
or _____ are actually produced.
In addition when dilute solutions of ____________________ are electrolyzed
REDUCTION: 2H2O + 2e- → H2(g) + 2OH- (10-7 M) E� = - 0.41 V
Fe2+ + 2e- → Fe E� = - 0.45 V
Cr3+ + 3e- → Cr E� = - 0.74 V
Zn2+ + 2e- → Zn E� = - 0.76 V
Strongest Oxidizing
Agent
we would expect that _____ is produced but in practice, ____________________ will be
produced.
Electrolysis of aqueous solutions containing Cl- or Br- will produce Cl2 or Br2 at the
anode.
Electrolysis of aqueous solutions containing Fe2+, Cr3+, or Zn2+ will produce Fe, Cr,
or Zn at the cathode.
Unit 5 Electrochemistry (S).doc 51
EXAMPLE 5.13 ELECTROLYSIS REACTIONS
Problem: What products are formed at the anode and cathode wand what is the overall reaction when a solution containing NiSO4(aq) is electrolyzed using inert electrodes? Determine the minimum voltage required.
Solution: First, the species present ______________________________ and the possible reactions are
REDUCTION
+ 4H+ + 4e- → H2SO3 + H2O E� = +0.17 V
+ 2e- → Ni E� = - 0.26 V
+ 2e- → H2 + 2OH- (10-7 M) E� = -0.41 V
OXIDATION
12 O2(g) + 2H+ (10-7 M) + 2e- ← E� = +0.82 V
______________________________________________________________
______________________________________________________________
so the half -reactions that occur are:
Anode: H2O → 12 O2(g) + 2H+ (10-7 M) + 2e- E� =
Cathode: Ni2+ + 2e- → Ni E� = - 0.26 V
H2O + Ni2+ → 12 O2(g) + 2H+ (10-7 M) + Ni E�CELL =
_____ is produced at the anode, _____ is produced at the cathode and a
voltage of at least __________ must be applied for the reaction to occur.
52 Chemistry 12
EXAMPLE 5.14 ELECTROLYSIS REACTIONS
Problem: What is the overall reaction which occurs when a 1 M solution of HCl(aq) is electrolyzed using carbon electrodes?
Solution: First, the species present H+, Cl -, H2O (acidic) and the possible reactions are
REDUCTION
+ 2e- → H2 E� = +0.00 V
+ 2e- → H2 + 2OH- (10-7 M) E� = -0.41 V
OXIDATION
Cl2(g) + 2e ← E� = +1.36 V
12 O2(g) + 2H+ (10-7 M) + 2e- ← E� = +0.82 V
-
Due to the high overpotential of H2O, the Cl- is oxidized so the two half reactions are
H+ + 2e- → H2 E� = + 0.00 V
2Cl- → Cl2(g) + 2e- E� =
H+ + 2Cl- → H2 + Cl2 E�CELL =
Anode
Cathode
_____ is produced at the anode, _____ is produced at the cathode and a
voltage of at least __________ must be applied for the reaction to occur.
Unit 5 Electrochemistry (S).doc 53
5. ELECTROPLATING
ELECTROPLATING is __________________________________________________
________________________________________________________________________
________________________________________________________________________
• The ____________________is made out of the material which will receive the metal
plating.
• The _________________________ ____________________ contains ions of the
metal which is to be “plated” onto the cathode.
• The _______________ may be made of the same metal which is to be “plated out”
onto the cathode. (This is normal but an inert electrode can also be used.)
54 Chemistry 12
EXAMPLE 5.15 ELECTROPLATING CELLS
Problem: Design a cell to electroplate a copper medallion with nickel metal. Include in the design: the ions present in the solution, the direction of ion flow, the substance used for the anode and cathode, and the direction of electron flow when the cell is connected to a DC power source.
Solution: Since nickel metal must be produce, the reduction reaction is:
The cell is setup as follows:
DC Power Source
• The medallion must have the Ni(s) plated onto it, so it must be the _______________.
• Make the _______________ out of Ni(s) so that the oxidation reaction
Ni(s) → Ni2+ + 2e-
will ensure a _________________________________________________
___________________________________________________________
• Since electrons always flow from ______________________________, connect DC power in such a way as to supply electrons to the cathode.
• The ions needed in solution will Ni2+ and some anion that does NOT form a precipitate with Ni2+. The NO3
- is always a good choice since all nitrates are soluble.
• As with spontaneous reactions, __________________________________
___________________________________________________________
___________________________________________________________
Unit 5 Electrochemistry (S).doc 55
6. ELECTROREFINING
ELECTROREFINING is _________________________________________________
________________________________________________________________________
________________________________________________________________________
DC Power Source
At the anode, the small amounts of Zn or Pb present is preferentially oxidized as it is
exposed at the surface. When any exposed Pb/Zn atoms have oxidized and gone into
solution as ions, only the Cu atoms are available to be oxidized. Any Au, Ag, or Pt
atoms present cannot be oxidized because the anode is mostly copper which is oxidized
in preference to Au, Ag, or Pt which simply drop off and accumulate on the bottom of the
cell. This “anode sludge” can be purified to obtain the valuable metals.