Dad

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Dad o the yearcircled the bowl
From a few years ago Guysyour son WILL do
something like this to you LadiesRemember to
forgive them (both) Yes, he was hiding the Dry
Erase marker behind his back
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Roadmap of stuff in Chapter 17
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Thermodynamics, Equilibrium, and
Electrochemistry A Summary
From any one of the three quantities Keq, ?G,
Ecell, we can determine the others.
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OxidationReductionThe Transfer of Electrons
Silver metal is formed, and the solution turns
blue from copper(II) ions formed.
Electrons from copper metal are transferred to
silver ions.
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A Qualitative Description of Voltaic Cells

• A voltaic cell uses a spontaneous redox reaction
  to produce electricity.
• A half-cell consists of an electrode (strip of
  metal or other conductor) immersed in a solution
  of ions.

This Zn2 becomes a Zn atom.
Both oxidation and reduction occur at the
electrode surface, and equilibrium is reached.
This Zn atom leaves the surface to become a Zn2
ion.
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Important Electrochemical Terms

• An electrochemical cell consists of two
  half-cells with the appropriate connections
  between electrodes and solutions.
• Two half-cells may be joined by a salt bridge
  that permits migration of ions, without
  completely mixing the solutions.
• The anode is the electrode at which oxidation
  occurs.
• The cathode is the electrode at which reduction
  occurs.
• In a voltaic cell, a spontaneous redox reaction
  occurs and current is generated.
• Cell potential (Ecell) is the potential
  difference in volts between anode and cathode.
• Ecell is the driving force that moves electrons
  and ions.

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A ZincCopper Voltaic Cell
Positive and negative ions move through the salt
bridge to equalize the charge.
the electrons produced move through the wire
to the Cu(s) electrode, where they are accepted
by Cu2 ions to form more Cu(s).
Zn(s) is oxidized to Zn2 ions, and
Reaction Zn(s) Cu2(aq) ? Cu(s) Zn2(aq)
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Cell Diagrams

• A cell diagram is shorthand for an
  electrochemical cell.
• The anode is placed on the left side of the
  diagram.
• The cathode is placed on the right side.
• A single vertical line ( ) represents a
  boundary between phases, such as between an
  electrode and a solution.
• A double vertical line ( ) represents a salt
  bridge or porous barrier separating two
  half-cells.

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• Describe the half-reactions and the overall
  reaction that occur in the voltaic cell
  represented by the cell diagram
• Pt(s) Fe2(aq), Fe3(aq) Cl(aq) Cl2(g)
  Pt(s)

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Standard Electrode Potentials

• Since an electrode represents only a
  half-reaction, it is not possible to measure the
  absolute potential of an electrode.
• The standard hydrogen electrode (SHE) provides a
  reference for measurement of other electrode
  potentials.

• The SHE is arbitrarily assigned a potential of
  0.000 V.

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Standard Electrode Potentials (contd)

• The standard electrode potential, E, is based on
  the tendency for reduction to occur at an
  electrode.
• E for the standard hydrogen electrode is
  arbitrarily assigned a value of 0.000 V.
• All other values of E are determined relative to
  the standard hydrogen electrode.
• The standard cell potential (Ecell) is the
  difference between E of the cathode and E of
  the anode.
• Ecell E(cathode) E(anode)

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Important Points About Electrode and Cell
Potentials
Electrode potentials and cell voltages are
intensive properties independent of the amount
of matter
Cell voltages can be ascribed to
oxidationreduction reactions without regard to
voltaic cells
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Cu2/Cu Electrode
EOS
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Zn2/Zn Electrode
EOS
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Observed Voltages
EOS
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Selected StandardElectrode Potentials at 25 oC
EOS
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• Determine E for the reduction half-reaction
• Ce4(aq) e ? Ce3(aq), given that the cell
  voltage for the voltaic cell
• Co(s) Co2(1 M) Ce4(1 M), Ce3(1 M)
  Pt(s)
• is Ecell 1.887 V.
• Balance the following oxidationreduction
  equation, and determine Ecell for the reaction.
• O2(g) H(aq) I(aq) ? H2O(l) I2(s)

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Criteria for Spontaneous Change
If Ecell is positive, the reaction in the forward
direction (from left to right) is spontaneous
If Ecell is negative, the reaction is
nonspontaneous If Ecell 0, the system is at
equilibrium
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Equilibrium Constantsfor Redox Reactions
DGo RTlnKeq n F Eocell
Eocell standard cell potential R is the gas
constant (8.3145 J mol1 K1) T is the Kelvin
temperature n is the number of moles of electrons
involved in the reaction F is the faraday constant
EOS
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• Where we left off or -, spontaneous or not
• Will copper metal displace silver ion from
  aqueous solution? That is, does the reaction
• Cu(s) 2 Ag(1 M) ? Cu2(1 M) 2 Ag(s)
• occur spontaneously from left to right?

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• Revisiting the copper/silver reaction
• Calculate the values of ?G and Keq at 25 C for
  the reaction
• Cu(s) 2 Ag(1 M) ? Cu2(1 M) 2 Ag(s)

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Using Standard Reduction Potentials
E -(-0.76 V)
E 0.34 V
E 1.10 V
E 0.80 V
E -0.34 V
E 0.46 V
Half-cell potentials are intensive properties.
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The Nernst Equation
DG DG RT ln Q
Using DG -nFE and DG -nFE
Nernst Equation
or
or
in volts, at 25C
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The Nernst Equation
Consider a galvanic cell that uses the reaction
What is the potential of a cell at 25 C that has
the following ion concentrations?
Fe2 0.20 M
Fe3 1.0 x 10-4 M
Cu2 0.25 M
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The Nernst Equation
Calculate E
E -0.34 V
E 0.77 V
Ecell -0.34 V 0.77 V 0.43 V
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The Nernst Equation
Calculate E
log
E E -
log
0.43 V -
E 0.25 V
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Standard Cell Potentials and Equilibrium Constants
DG -RT ln K
DG -nFE
and
Using
-nFE -RT ln K
log K
ln K

E
in volts, at 25C
log K
E
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Standard Cell Potentials and Equilibrium Constants
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Standard Cell Potentials and Equilibrium Constants
Three methods to determine equilibrium constants

1. K from concentration data

1. K from thermochemical data

1. K from electrochemical data

or
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Summary of Important Relationships
EOS
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Standard Reduction Potentials