Ch 19: Electrochemistry
Lecture Outline
Identify element being oxidized and element being reduced
1. Oxidation: Increasing oxidation number
2. Reduction: Decreasing oxidation number
3. Hierarchical guidelines
i. Free elements have an oxidation # of zero (e.g, Fe, H2, S8)
ii. The oxidation number must add up to the overall charge of the molecule or ion
iii. Assign oxidation numbers to all but one of the elements using the hierarchical guidelines. Find
the oxidation number of the final element mathematically.
1. Alkali metal ions (Group 1A) always have an oxidation number of +1. Alkaline earth
metals (Group 2A) are assigned an oxidation number of -1.
2. Fluorine, in a compound, is assigned -1
3. Hydrogen is usually +1 (less common oxidation number is -1)
4. Oxygen is usually assigned -2 (less common numbers include -1, peroxide)
4. OXIDIZED
i. Charge increases:
Example:
ii. Gains Oxygen and/or loses hydrogen (hydrocarbons): Example:
5. REDUCED
i. Charge decreases:
Example
ii. Gains hydrogen and/or loses oxygen (hydrocarbons)
6. OXIDIZING AGENT:
7. REDUCING AGENT:
Oxidation and Reduction: Balancing using half reaction method
-1-
Ch 19: Electrochemistry
Lecture Outline
Balance RedOx reactions by the half-reaction method
1. Assign oxidation numbers to all atoms, Identify what is being oxidized and what is being reduced.
2. Determine if the reaction is in acid or basic conditions (one extra step is required if the reaction occurs
in basic conditions)
3. Split the reaction into half-reactions.
a. Oxidation half reaction
b. Reduction half reaction
4. Balance, in order
a. All atoms except H and O
b. Balance the O atoms by adding H2O to the side of the equation that needs O
c. Balance H by adding H+ to the side of the equation that needs H.
d. Balance the electric charge by adding electrons (e-) to the more positive side
5. Combine the two half-reactions
a. Multiply each half-reaction by a factor so that each has the same number of electrons.
b. Add the two half-reactions (the electrons should cancel)
c. Simplify the reaction by canceling species that occur on both side of the reaction
d. Check that the reaction is balanced
6. If the reaction occurs in acidic conditions, you are finished. If it occurs in basic conditions, finish by
a. Adding one OH- to both sides of the reaction for each H+ (or H3O+)
b. When H+ and OH- occur on the same side, combine them to form H2O (or H3O+ + OH- makes two
waters)
c. Cancel water molecules that occur on both sides.
Oxidation and Reduction: Balancing using half reaction method
-2-
Ch 19: Electrochemistry
Lecture Outline
Example 1 : In an acidified solution, Dichromate ion (Cr2O72-) reacts with zinc metal. Products of the reaction include
zinc ions (Zn2+) and chromium(III) ions (Cr3+)
Write the balanced ionic equation for this reaction using the half−reaction method.
1.
Determine what is being oxidized and what is being reduced
a. Assign oxidation numbers
Cr2O72-
Zn2+
Zn
Cr3+
b. Element Reduced: _______________, element oxidized: __________________
Oxidizing Agent:
Reducing Agent
2. Reaction conditions: Acid or basic?
3. Oxidation half reaction:
Reduction half reaction:
4. Balance each half reaction
a. All atoms except H and O
b. Balance O (add H2O)
c. Balance the H (add H+)
d. Balance the charge (e-)
5. Combine the two half reactions.
a. Multiply (electrons must be equal)
b. Add
c. Simplify
d. check
6. Adjust for basic conditions if necessary
Oxidation and Reduction: Balancing using half reaction method
-3-
Ch 19: Electrochemistry
Lecture Outline
Example 2: Write a balanced net ionic reaction: In a basic solution, Lead(II) ion (Pb2+) react with hypochlorite ion (ClO-)
to give lead(IV) oxide (PbO2) and chloride ion (Cl-).
More examples:
Acidic: Zn(s) + NO3-(aq)  Zn2+(aq) + NH4+(aq)
Basic: H2O2 + ClO2  ClO2- + O2
Oxidation and Reduction: Balancing using half reaction method
-4-
Ch 19: Electrochemistry
Lecture Outline
VOLTAIC CELLS (aka galvanic cell) and Batteries
I.
A BATTERY
1. A package of one or more voltaic cells connected together to produce electric energy
III.
2. Word origin: A set or series of similar units as in a number of pieces of artillery used together
Voltaic Cells: Key ideas
1. Spontaneous __________________ reaction (ΔG<0) is the source of energy
2. Two half reactions are separated physically from each other
3. Connections are made two ways
i. Connection to the device needing power ( ____________ connection)
ii. Salt Bridge. Allows for ____________ to flow. (_______________ connection)
An atomic view of a voltaic cell:
IV.
ACTIVITY “Batteries”: How does a voltaic cell work. Objectives
II.
1. Identify the anode and the cathode in a drawing given either
The direction of the electron flow OR
The half-cell reactions OR
The activity series of metals
2. Describe the purpose of the salt bridge in a voltaic cell
Voltaic Cell (Galvanic Cell)
-5-
Batteries
How does a battery (voltaic cell) work?
Why?
When we use portable devices like MP3 players and cell phones we need a ready source of electricity to
provide a flow of electrons. Batteries are the common solution to this challenge. In a battery or voltaic
cell, oxidation and reduction reactions provide electrons which power our devices. In this activity we will
explore the chemistry of voltaic cells or batteries.
Model 1 – Voltaic Cell
Bulb not lit
Wire not
connected
Bulb lit
Wire
connected
Salt Bridge
–
Cl
Zn
K+
–
Cl
Cu
Salt Bridge
Zn
Cu
+
K
+
K
–
NO3
Cl
NO3–
Zn2+
Cl–
K+
NO3–
–
Zn2+ Cl
NO3–
NO3–
NO3–
Anode
NO3
Time Passes
–
NO3–
Zn2+
NO3–
–
Cu2+
NO3–
–
K+
Zn2+ NO3
Cu2+ NO –
3
NO3–
NO3–
2+
– Zn
Cl
NO3–
Cu2+
NO3–
K+
Cathode
Before the wire
is connected
A few minutes after the
wire was connected
1. Consider the reaction in Model 1. Notice that there are two zinc ions (Zn2+) in the beaker on the
left before the wire is connected. Explain why the number of nitrate ions (NO3–) present in that
beaker is correct.
2. Examine the system in Model 1 both before the wire is connected and after it is connected.
Identify two specific pieces of evidence that a chemical reaction has occurred as time passes with
the wire connected.
Batteries1
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3. Examine the diagram in Model 1.
a. Which piece of solid metal loses mass (gets smaller) as the reaction proceeds?
b. Does the number of zinc ions (Zn2+) in solution increase or decrease as the reaction proceeds?
c. Circle the half-reaction below that represents the change in the metal identified in part a.
Zn2+(aq) + 2e– → Zn(s) Zn(s) → Zn2+(aq) + 2e–
Cu2+(aq) + 2e– → Cu(s) Cu(s) → Cu2+(aq) + 2e–
d. Is the reaction circled in part c an oxidation or reduction reaction?
4. Examine the diagram in Model 1.
a. Which piece of solid metal is gaining mass as the reaction proceeds?
b. Where do those metal atoms come from? Explain.
c. Circle the half-reaction below that represents the change in the metal identified in part a.
Zn2+(aq) + 2e– → Zn(s) Zn(s) → Zn2+(aq) + 2e–
Cu2+(aq) + 2e– → Cu(s) Cu(s) → Cu2+(aq) + 2e–
d. Is the reaction circled in part c an oxidation or reduction reaction?
5. Electricity is the flow of electrons. Look back at Model 1. In which diagram can the electrons
flow through the wire, the one when the bulb is not lit or the one when the bulb is lit? (Circle
one.) Explain your answer.
6. Based on your answers in Questions 3 and 4, which piece of solid metal is giving up electrons,
and therefore losing them into the wire?
7. On the drawing of the voltaic cell in Model 1, draw an arrow to depict the direction that the
electrons are traveling through the wire.
2POGIL™ Activities for High School Chemistry
-7-
8. Consider the reactions occurring in Model 1.
a. What type of half-reaction (oxidation or reduction) is occurring at the piece of metal labeled
anode in Model 1?
b. What type of half-reaction (oxidation or reduction) is occurring at the piece of metal labeled
cathode in Model 1?
c. Explain how the phrase “an ox and a red cat” can help students remember the type of halfreaction that occurs at each electrode in an electrochemical cell.
9. Explain how the direction of electron flow in a voltaic cell is consistent with what you would
predict from the activity series?
10. Draw an unconnected voltaic cell similar to the one on the left side of Model 1 using iron and
silver as electrodes. The solutions should include silver ions (Ag+), iron ions (Fe3+), and nitrate
ions (NO3–). Use a salt bridge identical to that in Model 1.
11. Use the activity series of metals to determine which metal in the voltaic cell in Question 10
should be the anode. Explain your choice.
Batteries3
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12. Draw an arrow on your diagram in Question 10 to indicate the direction of electron flow
through the voltaic cell once the wire is connected.
13. Use Model 1 to complete the table.
Initially
After a Few
Minutes
Number of zinc ions (Zn2+) dissolved in solution
Number of copper ions (Cu2+) dissolved in solution
Number of nitrate ions (NO3–) dissolved in solution
Number of potassium ions (K+) dissolved in solution
Number of chloride ions (Cl–) dissolved in solution
14. Even though ions and electrons move around in a voltaic cell, the cell must stay electrically
neutral.
a. Explain how the anode half-cell in Model 1 remains electrically neutral (no charge) even
though zinc ions are being formed from neutral zinc metal. Refer to the table in Question 13
to support your answer.
b. Explain how the cathode half-cell in Model 1 remains electrically neutral (no charge) even
though copper ions are being removed from the solution. Refer to the table in Question 13 to
support your answer.
c. What is the role of the salt-bridge in a voltaic cell?
4POGIL™ Activities for High School Chemistry
-9-
Extension Questions
15. Work with your group to apply everything you have learned about batteries to label the voltaic
cell diagram below. Use the path and direction that the electrons are traveling to help you.
a. Label the following items on the diagram.
Site of oxidation
Site of reduction
Anode
Cathode
b. Complete the “few minutes after” drawing to show what ions would be in the beakers and the
salt bridge, and how the electrodes may have changed.
c. Write the oxidation and reduction half-reactions for the voltaic cell.
Wire not
connected
Salt Bridge
Salt Bridge
Cl–
Fe
K
K+
Cl
–
K+
Cl–
NO3–
–
NO3
NO3–
NO3–
Zn
Time Passes
NO3–
Fe2+
NO3–
Fe
Zn
+
NO3–
Fe2+
Zn2+
Zn2+ NO –
3
A few minutes after the
wire was connected
Before the wire
is connected
16. Propose a reason why it is necessary to place the metal electrodes in solutions of their own ions to
make a battery.
Batteries5
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Ch 19: Electrochemistry
Lecture Outline
Standard Electrode (Reduction) Potentials
Which direction will the electrons flow?
Voltaic Cell Notation
The oxidation half−cell, the ________, is written on the ______.
The reduction half−cell, the ________, is written on the ______.
Species oxidized I Oxidation product
II Species reduced I Reduction product
INTERPRET the following
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Write the reaction for the voltaic cell that is represented by the
following :
Al(s) | Al3+(aq) || Cu2+(aq) | Cu(s)
Sketch the voltaic cell represented by the above reaction. Label the
anode and cathode. Show the directions the electrons flow (external
connection) and the direction the ions flow (internal connection)
Voltaic Cell (Galvanic Cell)
- 11 -
Ch 19: Electrochemistry
Lecture Outline
Cell Potential (EMF: Electromotive Force)
MOVEMENT OF ELECTRONS THROUGH THE EXTERNAL
CIRCUIT:
Just as work is required to pump water from one point to
another, so work is required to move electrons.
Water flows from areas of high pressure to areas of low
pressure. Similarly, electrons flow from high electric
potential to low electric potential. Electric potential can be
thought of as electric pressure.
•
•
Spontaneous flow of electrons is due to the diff in
potential energy between two substance)
Cell potential (Ecell) is measured in VOLTS
(Joule/Coulomb)
Summary
Eo cell
ΔG
Q&K
Relationship
Reaction
Direction
Spontaneity
(as written)
The standard electrode potential, E°, is the electrode
potential when the concentrations of solutes are 1 M,
the gas pressures are 1 atm, and the temperature has a
specified value (usually 25°C). The superscript degree
sign (°) signifies standard−state conditions.
Cell Potential (EMF)
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Ch 19: Electrochemistry
Lecture Outline
Summary of electrochemistry units and constants
Joule (J): Unit of work
Work = -nFEcell (where n is the number of moles of
electrons involved in either half-cell reaction)
Coulomb (C) : Unit of charge
1 e- = 1.60218 X 10-19 C
Faraday Constant (F): The charge of one mole of electrons
9.6485 ×104 𝐶𝑜𝑢𝑙𝑜𝑚𝑏𝑠
1 𝑚𝑜𝑙𝑒 𝑜𝑓 𝑒 −
1F=
Volt (V): Energy absorbed (or evolved) when a charge
moves through an energy difference
𝐽
1𝑉 =
𝐶
Watt: Unit of Power
𝐽
1 W= 𝑠
Ampere: Unit of current
𝐶
1𝐴=
𝑆
The emf of a particular cell is 0.500 V. The cell reaction is
2Al(s) + 3Cu2+(aq)  2Al3+(aq) + 3Cu(s)
Calculate the maximum electrical work (J) of this cell.
Hint: To determine the value of n we need to examine
the half−reactions:
Electrochemistry units: Work done by a cell
- 13 -
Ch 19: Electrochemistry
Lecture Outline
Calculating Ecell from half reactions
Ecell = Eoxidation + Ereduction
Look back at the electrode potentials given in the Table
19.1. Are the electrode potentials given for oxidation or
reduction?
How can you determine the potentials for oxidation?
Example:
How does the reduction potential of Li+(aq) compare to
the oxidation potential of Li(s)? Write each reaction and
its cell potential.
Voltage is an Intensive Property
This means that if you double the coefficents for the
reaction the value of the standard potential
__________________________.
Give the standard potentials for the following
reactions:
Al(s) → Al3+(aq) +3e2Al(s) → 2Al3+(aq) +6e-
Sn2+(aq) +2e- → Sn(s)
3Sn2+(aq) +6e- → 3Sn(s)
Calculating E0cell
- 14 -
Ch 19: Electrochemistry
Lecture Outline
Let’s find Eºcell for the following cell:
Al(s) | Al3+(aq) || Cu2+(aq) | Cu(s)
A fuel cell is simply a voltaic cell that uses a
continuous supply of electrode materials to provide
a continuous supply of electrical energy. A fuel cell
employed by NASA on spacecraft uses hydrogen and
oxygen under basic conditions to produce electricity.
The water produced in this way can be used for
drinking. The net reaction is
2H2O(l) + 2e− ⇄ H2(g) + 2OH−(aq)
O2(g) + 2H2O(l) + 4e− ⇄ 4OH−(aq)
2H2(g) + O2(g)  2H2O(g)
Calculate the standard emf of the oxygen–hydrogen
fuel cell.
Calculating E0cell
- 15 -
E° = –0.83 V
E° = 0.40 V
Ch 19: Electrochemistry
Lecture Outline
Predicting the Direction of a Reaction: Standard Cell Potential (Eº) to Standard Free Energy (ΔGº)
Eo cell
ΔG
Q&K
Relationship
Reaction
Direction
Standard potentials
Cr2O72− + 14H+ + 6 e- ⇌ Cr3+ + 7H2O
MnO4− + 8H+ + 5 e- ⇌ Mn2+ + 4H2O
Will dichromate ion oxidize manganese(II) ion to
permanganate ion in acid solution under standard
conditions?
Spontaneity
(as written)
E°= 1.33 V
E° = 1.49 V
The free−energy change, DG, for a reaction equals the
maximum useful work of the reaction.
ΔG° = wmax = –nFE°
First we need determine the number of electrons
involved in the reaction.
From the cell potential, calculate the standard
free−energy change for the net reaction used in the
hydrogen–oxygen fuel cell:
2H2(g) + O2(g)  2H2O(l)
The cell potential is 1.23 V.
How does this compare with ΔGf° of H2O(l)?
(ΔGf° = -285.8 kJ/mol)
Calculating E0cell
- 16 -
Ch 19: Electrochemistry
Lecture Outline
A voltaic cell consists of one half−cell with Fe dipping into
an aqueous solution of 1.0 M FeCl2 and the other half−cell
with Ag dipping into an aqueous solution of 1.0 M AgNO3.
I Fe(s)
Ag (aq) I Ag(s)
Fe2+(aq)
+
E° = –0.41 V
E° = 0.80 V
Determine the reaction and calculate the cell potential
(ΔGf° Ag+(aq) = 77 kJ/mol, and Fe2+(aq)= –85
kJ/mol.)
Continued with same reaction:
Calculate ΔGrxn° from ΔGf°
Continued: Verify the relationship: ΔG°=-nFE°
Equilibrium Constants from Cell Potentials
nFE° = RT ln K
Rearrange, solve for E°
Calculating E0cell
- 17 -
Ch 19: Electrochemistry
Lecture Outline
Calculate the equilibrium constant K at 25°C for the
following reaction for the standard cell potential:
Pb2+(aq) + Fe(s) ⇄ Pb(s) + Fe2+(aq)
 Find the cell potential
 Now find K
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Ch 19: Electrochemistry
Lecture Outline
Dependence of Cell Potential on Concentration: Nernst
Equation
ΔG = ΔG° + RT ln Q
o
Ecell  Ecell
–
RT
ln Q
nF
This reaction is sometimes written as
0.02568
o
Ecell  Ecell
–
ln Q
n
Consider a voltaic cell,
Fe(s) | Fe2+(aq) || Cu2+(aq) | Cu(s), being run under
standard conditions.
a. Is ΔG° positive or negative for this process?
b. Change the concentrations from their standard
values in such a way that Ecell is reduced. Write
your answer using the shorthand notation.
What is the cell potential of the following voltaic cell at
25ºC?
Zn(s) | Zn2+(0.200 M) || Ag+(0.00200 M) |Ag (s)
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Ch 19: Electrochemistry
Lecture Outline
Determination of pH (ion selective electrodes)
Standard Hydrogen electrode:
Glass electrode || Test Solution
Ag(s) | Ag+(1 M) || H+ (test solution) | H2(1 atm) | Pt
What are the reactions?
Anode:
Cathode:
Overall:
Nernst Equation:
Expression for Q:
How is pH determined?
A pH meter is constructed using hydrogen gas bubbling over an inert platinum
electrode (the hydrogen electrode) at a pressure of 1.2 atm. The other electrode is
aluminum metal immersed in a 0.20 M Al3+ solution. What is the cell potential when
the hydrogen electrode is immersed in a sample of acid rain with a pH of 4.0 at 25°C?
Al(s) Al3+(0.20 M) || H+ (test solution) | H2(1.2 atm) | Pt
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Corrosion Control: Cathodic Protection
- 20 -
Ch 19: Electrochemistry
Lecture Outline
Common Voltaic Commercial Voltaic Cell
Zinc-Carbon Dry Cell (Leclanché )
Anode: Zn(s)  Zn2+(aq) + 2e−
Cathode:2NH4+(aq) + 2MnO2(s) + 2e−  Mn2O3(s) + H2O(l) + 2NH3(aq)
Alkaline Battery
By © Túrelio (via Wikimedia-Commons), 2009 /, CC BY-SA 3.0 de,
https://commons.wikimedia.org/w/index.php?curid=8286636
Zn(s) + 2OH−(aq) → ZnO(s) + H2O(l) + 2e− [E° = -1.28 V]
2MnO2(s) + H2O(l) + 2e− → Mn2O3(s) + 2OH−(aq) [E° = +0.15 V]
Overall reaction:
Zn(s) + 2MnO2(s) ⇌ ZnO(s) + Mn2O3(s) [e° = +1.43 V]
Recharable NiCD battery
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Corrosion Control: Cathodic Protection
- 21 -
Ch 19: Electrochemistry
Lecture Outline
During Discharge:
Eoxº=
The reactions at the nickel oxide electrode are:
Eredº=
The net reaction during discharge is
Lead Storage Cell:
Anode:
Pb(s) + HSO4−(aq)  PbSO4(s) + H+(aq) + 2e−
Cathode:
PbO2(s) + 3H+(aq) + HSO4−(aq) + 2e− PbSO4(s) + 2H2O(l)
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Corrosion Control: Cathodic Protection
- 22 -
Ch 19: Electrochemistry
Lecture Outline
Cathodic Protection: Corrosion Control:
Rusting occurs when iron comes in contact with iron:
Anode:
Fe(s)  Fe2+(aq) + 2e−
Cathode:
O2(g) + 2H2O(l) + 4e−  4OH−(aq)
Sacrificial Anode: A more active metal such as Zn or Mg, (More
easily oxidized) is attached as
protection.
Examples: Compare oxidation
potentials of Fe and Zn
Zn (s)  Zn2+
Fe(s)  Fe2+
Applications of Cathodic Protection:
Galvanized Nails
Hull of Ship
Water Heaters
Pipelines
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Corrosion Control: Cathodic Protection
- 23 -
Ch 19: Electrochemistry
Lecture Outline
Electrolytic Cell
An electrolytic cell is an electrochemical cell in which electric current _______________ an otherwise
__________________ reaction. This process is called ________________________________.
(See “Electrolytic Cells”)
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Corrosion Control: Cathodic Protection
- 24 -
Name
Chem 163 Section: ______ Team Number: ______
ALE 27. Electrolytic Cells
(Reference: 21.7 Silberberg 5th edition)
If an element doesn’t occur naturally, where does a sample of it come from?
The Model: A Battery Used as an “Electron Pump”
If a battery is used to “pump” electrons from one electrode to
a second electrode, it is possible to force a non-spontaneous
reaction to occur. Let us at the moment consider an
electrolytic cell that has inert electrodes. An inert electrode
is made of some electrically-conducting substance, which is
resistive to being oxidized or reduced. Graphite or an
unreactive metal, like gold or platinum, may be used. In the
cell to the right, the electrodes are submerged in an aqueous
solution of some soluble salt, a strong electrolyte. (While the
salt is represented generically as “MX”, nothing is implied
about the charges of the cation and anion—their charges do
not have to be 1:1.)
Key Questions
1. As the battery pumps electrons from the electrode on the left to the electrode on the right,
what charge begins to accumulate on each electrode? (Circle your choices.)
Electrode on left: positive or negative
Electrode on right: positive or negative
2. What species flows to the electrode on the left? to the electrode on the right? Hints: Do like
charges attract or repel? Do opposite charges attract or repel? (Circle your choices.)
Electrode on left: anions or cations
Electrode on right: anions or cations
3. In each of the following half-reactions, are electrons reactants or products?
Oxidation: e- is reactant or product
Reduction: e- is reactant or product
4. In the electrolytic cell, which species are likely to be oxidized? to be reduced? (Circle all
choices that apply.)
Species likely to collide with anode and be oxidized:
anions
cations
water
Species likely to collide with cathode and be reduced: anions
cations
water
Page 1 of 7
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5
a. Now consider the electrolysis of Na2SO4(aq) using inert electrodes exposed to the air.
At the anode, either sulfate or water may oxidize:
2 SO42-(aq)

S2O82-(aq) + 2 e-
E oxo  -2.00 V
2 H2O(l)

O2(g) + 4 H+(aq) + 4 e-
E oxo  -1.23 V
Which species would you predict is the one that is oxidized at the anode? Briefly explain
why.
b. When Na2SO4(aq) is electrolyzed, either sodium cations or water may be reduced:
Na+(aq) + e- 
Na(s)
o
E red
 -2.71 V
2 H2O(l) + 2 e- 
H2(g) + 2 OH-(aq)
o
E red
 -0.83 V
Which species would you predict is the one that is reduced at the cathode? Briefly
explain why.
c. Suppose Na2SO4(aq) is
electrolyzed using inert
electrodes which are exposed to
the air. On the following
schematic of the electrolytic
cell…
i. draw what is being consumed
and what is being produced
at each of the electrodes.
ii. draw arrows indicating the
flow of each reactant and
product.
d. Why cannot pure water be electrolyzed? (i.e., Why must sodium sulfate or some similar
electrolyte be dissolved in the water? Hint: Consider the flow of ions in the aqueous
solution between the electrodes.)
Page 2 of 7
- 26 -
6. Let us now consider the electrolysis of CuSO4(aq) using inert electrodes exposed to the air.
The possible processes at the electrodes are:
at the anode
2 SO42-(aq)  S2O82-(aq) + 2 e-
at the cathode
E oxo  -2.00 V
2 H2O(l)  O2(g) + 4 H+(aq) + 4 e- E ox  -1.23 V
o
Cu2+(aq) + 2 e-  Cu(s)
o
E red
 0.34 V
2 H2O(l) + 2 e-  H2(g) + 2 OH-(aq) E red  -0.83 V
o
a. Circle the half-reaction that occurs at the anode and the reaction at the cathode.
b. Write the overall cell reaction and determine the absolute minimum voltage that the
battery must have to force this nonspontaneous reaction to occur. (I’ve said “absolute
minimum,” because ordinarily an additional voltage known as an overvoltage (a.k.a.
over potential or junction potential) must be supplied to overcome the slow reaction
kinetics due to the high activation energy required for gases to form at the electrode—
overvoltages vary, but are about 0.4 to 0.6 V for O2 and H2.

Overall reaction:

o
Minimum E cell
=
o
o
c. If overvoltage is considered, what’s the approximate E cell
? E cell
= __________________
The Model: Electrolysis of Molten Salt
If we want to avoid the possibility of more than one half-reaction
taking place at each electrode, the best way to do that is to limit
which species can interact with the anode and cathode. Suppose
instead of an aqueous solution of a salt (again, generically
represented as “MX”) we electrolyze the molten salt [i.e.,
“MX(l)”]. Then it is possible that both the cations of the salt
would be reduced to the elemental state at the cathode and the
anions of the salt would be oxidized to the elemental state at the
anode.
Key Questions
7. Why must a salt be in the liquid state if it is to be electrolyzed?
8
a. What is an advantage of electrolyzing a molten salt over electrolyzing an aqueous
solution of that salt?
b. What is a procedural challenge to electrolyzing a molten salt?
Page 3 of 7
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9. Elemental magnesium does not occur freely in Nature, but Mg2+ is the third most abundant
ion in sea water. At a concentration of 0.056 M, there is a virtually inexhaustible supply of
magnesium in our oceans! Once magnesium chloride salt is isolated from sea water, it can
be electrolyzed to yield magnesium metal. The light-weight magnesium metal is moderately
strong, so it is a useful metal for construction (e.g., of airplanes). The combustion of
magnesium metal produces an intense light, so it is also used in pyrotechnics.
a. Write the oxidation and reduction half-reactions that occur at the anode and cathode,
respectively, when MgCl2(l) is electrolyzed using inert electrodes.
Half reaction at anode:
Half reaction at cathode:
b. Elemental magnesium
metal, Mg, and chlorine
gas, Cl2, react explosively
with each other! So when
MgCl2(l) is electrolyzed,
the products must be
prevented from
recombining. Use the
design of the electrolytic
cell to the right and the
following physical data to
explain how these
elements may be collected
when molten magnesium
chloride, MgCl2(l), is
electrolyzed using inert
electrodes.
Mg
MgCl2
Tf
d
(C) (g/mL)
649
1.7
714
2.3
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The Model: Electroplating
Electroplating is the process of using electrical
current to reduce cations of a desired material from a
solution and coat a conductive object with a thin
layer of the material (e.g. such as a metal).
Electroplating is primarily used for depositing a
layer of material to give it a desired property (e.g.,
abrasion and wear resistance, corrosion protection,
aesthetic qualities, etc.) to a surface that otherwise
lacks that property.
Another application uses electroplating to build up
thickness on undersized parts as seen in the
electrolytic cell to the right where an external power
source is used to increase the thickness of the tin
cathode.
Key Questions
10. The possible processes at the electrodes are:
at the cathode
Sn2+(aq) + 2e-  Sn(s)
at the anode
o
E red
 -0.14 V
2 H2O(l) + 2 e-  H2(g) + 2 OH-(aq) E red  -0.83 V
o
Cu(s)  Cu2+(aq) + 2 e-
E oxo  -0.34 V
2 H2O(l)  O2(g) + 4 H+(aq) + 4 e- E ox  -1.23 V
o
a. Circle the half-reaction that occurs at the cathode and the half-reaction at the anode if the
voltage is kept “low.” (Water will also react if the voltage is high enough.)
b. Write the overall cell reaction and determine minimum voltage that the battery must have
to force this nonspontaneous reaction to occur. (Do not consider overvoltage since gases
are not involved.)
Overall reaction:
o
Minimum E cell
=
c. If water were to react at each electrode, what is the minimum voltage the battery must
have for water to under go hydrolysis and produce hydrogen gas at the cathode and
oxygen gas at the anode? (Since gases are involved remember to consider overvoltage)
What is the net reaction involving water at each electrode?
Overall reaction:
o
 the hydrolysis of water =
Minimum E cell
for
d. What is the minimum voltage (remember to consider overvoltage) the battery must have
for both hydrogen gas to form (but no O2 at the anode) and tin to plate on the cathode?
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The Model: Calculating the Mass of Metal Electroplated
A common method of electroplating is achieved using the object to be electroplated as the
cathode (e.g. the spoon in the figure below) and use as the anode the metal to be plated (e.g.
silver) both submerged in an aqueous solution containing the ions to be plated (e.g. AgNO3(aq)).
In the figure below, silver anode is oxidized (Ag(s)  Ag+(aq) + e-) and at the cathode Ag+ is
reduced (Ag+(aq) + e-  Ag(s)). Thus there is no net reaction. The overall process consists of
simply moving silver atoms from one electrode to the other. Since this is not a spontaneous
process, an electrical current is needed to drive the process. The mass in grams of silver
electroplated on the spoon is calculated using the following relationships:
The approximate mass of silver electroplated can be calculated
from the electrical current in amps and the amount of time the
current ran:
Current: 7.50 amps at a minimum of 0.80 V
Time: 30.0 minutes
Useful Info. Needed to Calculate the Mass of Ag Electroplated:




1 ampere (A) = 1 Coulomb per sec: 1 A = 1 C / s
Charge of 1 mol of electrons =96,485 C = 1 Faraday
1 F = 96,485 C / 1 mol e+
o
 0.80 V
Ag (aq) + e-  Ag (s) E red
o Hence, 1 mol of electrons will plate 1 mol of Ag
Molar Mass of Ag: 107.9 g/mol
Key Question
11. If the electrolytic cell above were run at 7.50 A for 30.0 minutes, use dimensional analysis to
show that the maximum mass of silver that can be electroplated is 15.1 g.
12. How long will it take (in hours) to electroplate 40.0 g of Chromium onto a small automotive
accessory from an aqueous Cr(NO3)3 solution using a current of 12 amps?
Cr = 51.0 g/mol. Use dimensional analysis to show your work.
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Exercises
13. Zinc plating (galvanizing) is an important means of corrosion protection. Although the
process is done customarily by dipping the object into molten zinc, the metal can also be
electroplated from an aqueous Zn2+ solution. How many grams of zinc can be deposited on a
steel tank from a ZnSO4 solution when a 0.755 amp current flows for 2.00 days? Use
dimensional analysis to show your work.
14. Electrolysis of molten NaCl is the major means of producing sodium metal. If 215 g of Na
metal forms… Use dimensional analysis to show your work and circle your answers.
a. How many moles of electrons are required?
b. How many coulombs are required?
c. How many amps are required to produce this amount in 9.5 hr?
15. In the electrolysis of a molten mixture of CsBr and SrCl2, use your knowledge of periodic
trends in ionization energy and electron affinity to identify the product that forms at the
negative electrode (cathode) and at the positive electrode (anode). Explain your reasoning.
Substance formed at the cathode: ______ Reasoning:
Substance formed at the anode: ______ Reasoning:
Page 7 of 7
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Ch 19: Electrochemistry
Lecture Outline
Voltaic Cells
1. Use the words anode and cathode to complete the sentences
 Oxidations occurs at the ____________ and reduction occurs and the ____________
 Electrons flow from the ______________ to the ________________
 Cations (from the salt bridge) flow to the _______ cell and anions flow to the _____________
 The electrode at the _________ is positive and the electrode at the ___________ is negative
 If both electrodes are active metals, the mass of the __________ increases but the mass of the
______________ decreases
 The concentration of metal ion at the ______ increases and the concentration of metal ion at the
_________ decreases
2. Practice a labeling exercise:
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Study Guide/Practice Questions (Midterm #3)
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Ch 19: Electrochemistry
Lecture Outline
3. Given a cell reaction, identify the reactions that happen at the cathode and anode.
Example: The reaction for a lead storage battery is
Pb(s) + PbO2(s) + 2 H+(aq) + 2 HSO4-(aq) 2 PbSO4(s) + 2 H2O(l)


What is the cathode and anode reactions?
What is oxidized, what is reduced?
4. Given two reduction potentials, decide what reaction occurs at the anode and cathode, and calculate the overall
potential of the cell. Also calculate ΔGº and the equilibrium constant K
Ni2+(aq) + 2e-  Ni(s) Eºred = -0.23V
Fe3+(aq) + 3e-  Fe(s) Eºred = -0.04V
5. Interpret short hand cell notation. Translate the following notations
 Fe(s) I Fe2+(aq) I I Ag+(aq) I Ag(s)
 Pt I H2(g) (1atm) I H+ (aq) (0.05 M) I Br2(l) I Br- (aq)(2.5M) I Pt
6. Non- standard condition cells
 Calculate the value of “Q” for the cell:
Pt I H2(g) (1atm) I H+ (aq) (0.05 M) II Br2(l) I Br- (aq)(2.5M) I Pt
 Calculate the potential for the above cell at 25 ºC
 What is the nickel(II) ion concentration in the voltaic cell if the cell potential is 0.34 V at 25 ºC Zn(s) I
Zn2+(aq)(1.00M) I I Ni2+(aq) (? M)I Ag(s)
Electrolytic Cell
7. Use the words anode and cathode to complete the sentences
 Oxidations occurs at the ____________ and reduction occurs and the ____________
 Electrons flow from the ______________ to the ________________
 Cations (from the salt bridge) flow to the _______ cell and anions flow to the _____________
 The electrode at the _________ is positive and the electrode at the ___________ is negative
 If both electrodes are active metals, the mass of the __________ increases but the mass of the
______________ decreases
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Study Guide/Practice Questions (Midterm #3)
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Ch 19: Electrochemistry
Lecture Outline
8. Electrolysis of water
Decide what is likely to happen during electrolysis of aqueous solutions. You must consider that water might
undergo oxidation or reduction.




What is the balanced half reaction for the reduction of water?
What is the balanced half reaction for the oxidation of water?
What are the anode and cathode reactions for the electrolysis of NaCl(l) ?
What are the anode and cathode reactions for the electrolysis of NaCl(aq)?
Stoichiometry Of Electrolysis
9. Important relationships to know
 What is a Volt?
 What is an Ampere?
 What is a Faraday?
10. A constant electric current deposits 365 mg of copper in 216 min from an aqueous copper(II) nitrate solution. What
is the current?
Spontaneous vs Non-spontaneous
11. Know the meaning of the sign and or magnitude of E, ΔG, and K.
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Study Guide/Practice Questions (Midterm #3)
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