Please Sit with Your Group

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Please Sit with Your Group

• Please be sure each member of your team has a
  copy of
• Electrochemical Cells Lecture Notes
• Electrochemical Cells Problem Set
• Todays reporter is the person whos box number
  has the most odd digits.
• Next reading assignment
• Zumdahl Chapter 11.5

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Electrochemical Cells

• Edward A. Mottel
• Integrated, First-Year Curriculum in Science,
  Engineering and Mathematics

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Electrochemical Cells

• Voltaic and electrolytic cells
• Electrochemical cell structure
• Types of electrochemical cells

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Electrochemical Cell

• A physical arrangement involving
• an oxidation reaction
• a reduction reaction

• The potential of the cell can be
• spontaneous
• Voltaic cell (also called a Galvanic cell)
• non-spontaneous
• Electrolytic cell

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Voltaic Cell

• An electrochemical cell which
• spontaneously generates a positive electrical
  potential
• can be used for useful work
• has Ecell gt 0 as constructed
• Example
• A discharging battery
• rechargeable or non-rechargeable

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Electrolytic Cell

• An electrochemical cell which
• requires an external energy source to force the
  cell in a non-spontaneous direction.
• has Ecell lt 0 as constructed.
• Examples
• A battery being recharged.
• A piece of metal being electroplated.

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Electrochemical Cell Structure
Cathode Electrode
Anode Electrode
Salt Bridge
Cathode Solution
Anode Solution
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Electrochemical Cell Structure
Cathode Electrode
Anode Electrode
Salt Bridge
Cathode Solution
Anode Solution
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Electrochemical Cell Structure
Cathode Electrode
Anode Electrode
Salt Bridge
Cathode Solution
Anode Solution
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Electrochemical Cell Structure
2.002 V
Cathode Electrode
Anode Electrode
Salt Bridge
Cathode Solution
Anode Solution
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Electrochemical Cell Structure

• Half-cell reactions
• Electrodes
• Electron flow
• Ion flow
• Shorthand notation

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Half-Cell Reactions

• Each electrochemical cell involves both an
  oxidation reaction and a reduction reaction.
• The oxidation cell and the reduction cell are
  referred to as half-cells.

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Anode Reaction
2.002 V
Cathode Electrode
Anode Electrode
Salt Bridge
Cathode Solution
Anode Solution
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AnodeThe electrode at which oxidation occurs
2.002 V
Cathode Electrode
Anode Electrode
Salt Bridge
Cathode Solution
Anode Solution
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Anode of a Voltaic Cell is Negative
2.002 V
Cathode Electrode
Anode Electrode
e
Salt Bridge
-
Al
Cathode Solution
Anode Solution
because electrons are released
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2.002 V
Cathode Electrode
Anode Electrode
Salt Bridge
-
Cathode Solution
Anode Solution
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Cathode Reaction
2.002 V
Cathode Electrode
Anode Electrode
Salt Bridge
-
Cathode Solution
Anode Solution
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CathodeThe electrode at which reduction occurs
2.002 V
Cathode Electrode
Anode Electrode
Salt Bridge
-
Cathode Solution
Anode Solution
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Cathode of a Voltaic Cell is Positive
2.002 V
Cathode Electrode
Anode Electrode
Salt Bridge
-

Cathode Solution
Anode Solution
because electrons are attracted and consumed
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2.002 V
Cathode Electrode
Anode Electrode
Salt Bridge
-

Cathode Solution
Anode Solution
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Electrons are transferredthrough a wire from
anode to cathode
2.002 V
Cathode Electrode
Anode Electrode
Salt Bridge
-

Cathode Solution
Anode Solution
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Electron Current Flowmay be used to perform
useful work
Cathode Electrode
Anode Electrode
Salt Bridge
-

Cathode Solution
Anode Solution
Electrical connection is made at the
electrodes, the site at which oxidation and
reduction occurs.
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Keeping It Straight
Anode
Cathode
Oxidation
Reduction
Electrons are attracted and consumed
Electrons are released
In a voltaic cell it is the negative electrode
In a voltaic cell it is the positive electrode
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Electron Flow

• Electrons are transferred through a wire from the
  anode to the cathode.

Ion Flow

• Anions are attracted to the anode and cations
  migrate away from anode.

Salt Bridge

• The salt bridge contains an ionic compound such
  as KNO3 or NaCl dissolved in a gel such as
  agar-agar.

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Draw a Diagram
indicate what is happening to all the charged
species in the anode cell.
List charged species
Show their location and their motion
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Anode Cell
Show the motion of all the charged species
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Ion Flow

• Cations are attracted to the cathode and anions
  migrate away from cathode.

Draw a diagram indicating what is happening to
all the charged species in the cathode cell.
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Cathode Cell
Identify the main species
Show the motion of all the charged species
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Salt Bridge

• A salt bridge may be used to physically separate
  ions in one half-cell from ions in the other
  half-cell.

Draw a diagram indicating what is happening to
all the charged species in the salt bridge.
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Salt Bridge
NO3
Al3
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Shorthand Line Notation
anode anode solution cathode solution
cathode

• Al(s) Al3 (1.00 M) Cu2 (1.00 M) Cu(s)

H2(g, 1 atm), Pt(s) H (1 M) Cl (1 M)
Cl2(g, 1 atm), C(gr)
Why is a graphite or a platinum electrode needed?
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Types of Electrochemical Cells

• Concentration Cell
• Standard Redox Cell
• Non-standard (Combination) Redox Cell

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Concentration Cell

• The oxidation and reduction reactions are
  identically reverse of each other.
• The observed cell potential is due solely to
  differences in concentrations of the solutions
  involved.
• Low potentials generated (mV)

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Concentration Cell

• Example
• Zn(s) Zn2 (0.23 M) Zn2 (1.00 M) Zn(s)

Write the oxidation and reduction
half-cell reactions taking place in this cell.
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Concentration Cell

• Example
• Zn(s) Zn2 (0.23 M) Zn2 (1.00 M) Zn(s)

Write the Q term for this cell.
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Concentration Cell

• Example
• Zn(s) Zn2 (0.23 M) Zn2 (1.00 M) Zn(s)

Determine the standard cell potential for this
cell.
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Concentration Cell

• Ecell 0.00 V
• Low potentials generated (mV)

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Standard Redox Cell

• The oxidation and reduction reactions are
  different.
• Concentrations of solutions are 1 M and reactant
  gas pressures are 1 atm.
• The observed cell potential is due to the
  differences in the activity of the reactants.

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Standard Redox Cell

• Example
• Ni(s) Ni2 (1.00 M) Ag (1.00 M) Ag(s)

Write the oxidation and reduction
half-cell reactions taking place in this cell.
Write the Q term for this cell.
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Standard Redox Cell

• Example
• Ni(s) Ni2 (1.00 M) Ag (1.00 M) Ag(s)

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Why is this called a standard redox cell?
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Standard Redox Cell

• Example
• Ni(s) Ni2 (1.00 M) Ag (1.00 M) Ag(s)

Determine the standard cell potential for this
cell.
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Standard Redox Cell

• Ecell ¹ 0.00 V
• Potentials (voltage) generated can be quite high

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Non-standard (Combination) Redox Cell

• The oxidation and reduction reactions are
  different.
• The solution concentrations are not 1 M.
• Gas pressures are not 1 atm.

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Non-standard (Combination) Redox Cell

• Example
• Mn(s) Mn2 (1.00 M) Pb2 (0.23 M) Pb(s)

Write the oxidation and reduction
half-cell reactions taking place in this cell.
Write the Q term for this cell.
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Non-standard (Combination) Redox Cell

• Example
• Mn(s) Mn2 (1.00 M) Pb2 (0.23 M) Pb(s)

4.3
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Non-standard (Combination) Redox Cell

• Example
• Mn(s) Mn2 (1.00 M) Pb2 (0.23 M) Pb(s)

Determine the standard cell potential for this
cell.
Why is this called a non-standard redox cell?
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Non-standard (Combination) Redox Cell

• The majority of the observed cell potential is
  due to the differences in the activity of the
  reactants, modified slightly by non-standard
  conditions.
• Ecell ¹ 0.00 V
• Potentials generated can be quite high (V)

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Electrode Materials

• Inert electrodes can or must be used in some
  instances.
• The reactant or product is a gas or liquid.
• The reactant and product of a half-cell are
  soluble.
• The product is being plated out onto an inert
  electrode.

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Inert Electrodes

• Examples
• Pt(s) Cr2 (1.00 M), Cr3 (1.00 M)
  Cu2 (1.00 M) Au(s)
• Co(s) Co2 (0.789 M)
  Hg2 (0.50 M) Hg(l),Pt(s)
• H2(g, 30 atm), C (gr) KOH (0.789 M)
  KOH (0.789 M) O2(g, 20 atm), C (gr)

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Pt(s) Cr2(1.0 M), Cr3(1.0 M) Cu2 (1.0 M)
Au(s)

• Draw a beaker diagram for this cell.
• Identify what is being oxidized and what is being
  reduced.
• Indicate the flow of all cations, anions and
  electrons in your diagram.
• What is the standard cell potential?
• What is the Q term?

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