Chapter Eighteen

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Chapter Eighteen

• Electrochemistry

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Oxidation-Reduction Transfer Of Electrons
Wait
Cu metal and Ag ions
Ag metal and Cu2 ions
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Concepts In Oxidation-Reduction

• Oxidation-reduction reactions involve a transfer
  of electrons from one species to another.
• Oxidation loss of e- or oxidation increases
• Reduction gain of e- or oxidation decreases
• The species that loses electrons is oxidized and
  is an reducing agent.
• The species that gains electrons is reduced and
  is an oxidizing agent.
• Review Sections 4.4 4.5 for details of
  oxidation-reduction reactions.

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Concepts In Oxidation-Reduction
Ag e ? Ag reduction of silver silver
is the oxidizing agent Cu ? Cu2 2e
oxidation of copper copper is the reducing
agent
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However, we never see electrons in the final
balanced chemical equation 2 Ag Cu ? 2
Ag Cu2 Can only see electrons when the
reaction is broken down into half
reactions Ag e? Ag Cu ? Cu2
2e Number of electrons produced and consumed in
half reactions must be balanced so that they
cancel out in the final chemical equation.
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Half-Reactions

• In any oxidation-reduction reaction, there are
  two half-reactions
• Ag e ? Ag reduction half-reaction
• Cu ? Cu2 2e oxidation
  half-reaction

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Half-Reaction Method OfBalancing Redox Equations

• Separate an oxidation-reduction equation into two
  half-equations, one oxidation and one for
  reduction.
• Balance the atoms and the electric charge in each
  half-reaction. Electrons appear on the left in
  the reduction half-equation and on the right in
  the oxidation half-equation.
• Adjust the coefficients in the half-equations so
  that the same number of electrons appears in each
  half-equation.
• Add together the two adjusted half-equations to
  obtain an overall oxidation-reduction equation.

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Redox Reactions In Acidic Solution

• In these reactions, it will be necessary to add
  molecules of water to the right and protons (H)
  on the left in the reduction step.
• In some of the oxidation reactions, it will be
  necessary to add molecules of water to the left
  and protons on the right.
• After these equations have been mass balanced,
  they will need to be charge balanced by adding
  electrons
• Simplify the overall equation, if necessary, by
  removing redundant species (species which appear
  on both sides of the equation).

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Redox Reactions In Acidic Solution

• Assign oxidation numbers to each atom and
  identify the species undergoing oxidation and
  reduction. Write two skeleton half-reactions.
• oxidation increase gt oxidation oxidation
  decrease gt reduction
• (2) Balance the numbers of atoms in each
  half-reaction.
• (a) Balance all atoms except O and H.
• (b) Balance O atoms by adding H2O molecules.
• (c) Balance H atoms by adding H ions.
• Balance electric charge by adding electrons to
  more positive side.
• Combine two half-reactions to obtain balanced
  equation.
• (a) Multiply each half-reaction by factors that
  make the number of
• electrons in two half-reactions equal.
• (b) Combine the adjusted half-equations into an
  overall equation.
• Simplify the balanced equation if necessary.
• Verify that the equation is balanced.

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Redox Reactions In Acidic Solution

• Balance the following redox equation in acidic
  solution
• Fe2 (aq) Cr2O72- (aq) ? Fe3 (aq) Cr3 (aq)

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Redox Reactions In Basic Solution

• For a reaction in basic solution, OH- should
  appear instead of H in the balanced equation.
• One method commonly used to balance such
  equations is to balance the equation as if the
  reaction occurs in acidic solution.
• Then, to each side of the net equation, add a
  number of OH- ions equal to the number of H
  appearing in the equation.
• As a result, one side of the equation will have
  H and OH- ions in equal number those can be
  combined into water.

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Redox Reactions In Basic Solution

• Steps (1) (5) same as in acidic solution.
• Convert from acidic to basic solution.
• (a) Add OH- to each side of the equation.
• (b) Combine H and OH- to H2O.
• (c) Simplify the equation if necessary.
• Verify
• Balance the following redox equation in basic
  solution
• S2- (aq) I2 (aq) ? SO42- (aq) I- (aq)

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Electrochemical Cells

• An electrochemical cell is a system consisting of
  electrodes that dip into an electrolyte and in
  which a chemical reaction either uses or
  generates an electric current.
• A voltaic cell is an electrochemical cell in
  which electric current is generated from a
  spontaneous redox reaction.

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A Zinc-Copper Voltaic Cell
Zn electrode Anode (-) Oxidation Zn ? Zn2 2 e-
Cu electrode Cathode () Reduction Cu2 2 e- ?
Cu
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Qualitative Description Of Voltaic Cells

• A half-cell consists of an electrode immersed in
  a solution of ions.
• The solutions in the two half-cells are joined by
  a salt bridge. This bridge keeps to the two
  half-cells in contact with one another so that
  there can be a flow of electrons.
• Metal wires connect the electrodes to the
  terminals of an electric meter called a
  voltmeter.
• The meter indicates that electrons flow
  continuously through the system.

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Electrode Equilibrium
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Important Electrochemical Terms

• Electrode is the metal strip used in
  electrochemical experiment.
• The anode is the electrode at which oxidation
  occurs and the cathode is the electrode at which
  reduction occurs.
• The cell potential (Ecell) is the potential
  difference that is the driving force that propels
  electrons from the anode to the cathode.
• Cell reaction is the redox reaction that occurs
  in a voltaic cell.

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Cell Diagrams

• Place the anode on the left side of the diagram.
• Place the cathode on the right side of the
  diagram.
• Use a single vertical line ( ) to represent the
  boundary between different phases, such as
  between an electrode and a solution.
• Use a double vertical line ( ) to represent a
  salt bridge or porous barrier separating two
  half-cells.
• When the half-reaction involves a gas, an inert
  material such as platinum serves as an electrode
  surface on which the half-reaction occurs.

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Cell Diagram
(see p. 775 text)
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Cell Diagrams Examples

• Write the cell diagram for a voltaic cell with
  the
• following redox reaction.
• 2 Al (s) 6 H (aq) ? 2 Al3 (aq) 3 H2 (g)
• Write cell reactions for the following voltaic
  cells.
• Cd(s) Cd2 (aq) Ni2 (aq) Ni (s)
• Zn (s) Zn2 (aq) Fe3 (aq), Fe2 (aq) Pt

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Standard Electrode Potentials

• A standard electrode potential, Eo, is based on
  the tendency for reduction to occur at the
  electrode.
• The cell voltage, called the standard cell
  potential (Eocell), is the difference between the
  standard potential of the cathode and that of the
  anode.
• Eocell Eo (cathode) Eo (anode)

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Standard Hydrogen Electrode

• In the standard hydrogen electrode, hydrogen gas
  at exactly 1 bar pressure is bubbled over an
  inert platinum electrode and into an aqueous
  solution with the concentration adjusted so that
  the activity of H3O is exactly equal to one.
• 2 H (a1) 2 e- ?H2 (g, 1 atm)
• E0 0 V

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Measuring The Standard PotentialOf The Cu2 / Cu
Electrode
Pt H2(g, 1atm) H(1M) Cu2(1M)
Cu(s) anode cathode Eocell Eo(Cu2/Cu)
Eo(H/H2) Eo(Cu2/Cu) 0.000V 0.340
V Eo(Cu2/Cu) 0.340 V
Cu2 ions are more readily reduced to Cu (s) than
H ions are reduced to H2.
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Measuring The Standard PotentialOf The Zn2 / Zn
Electrode
Pt H2(g, 1atm) H(1M) Zn2(1M)
Zn(s) anode cathode Eocell Eo(Zn2/Zn)
Eo(H/H2) Eo(Zn2/Zn) 0.000V - 0.763
V Eo(Zn2/Zn) - 0.763 V
Zn2 ions are less readily reduced to Zn (s) than
H ions are reduced to H2.
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Standard Electrode Potentials
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Important Points About Electrode And Cell
Potentials

• Electrode potentials and cell voltages are
  intensive properties. Their magnitudes are fixed
  once the particular species and their
  concentrations are specified. The magnitudes do
  not depend on the total amounts of the species
  present, for example, not on the size of a
  half-cell or voltaic cell.
• Cell voltages can be ascribed to
  oxidation-reduction reactions without regard to
  voltaic cells. Specifically, we can calculate
  Eocell from the equation for a cell reaction
  without writing a cell diagram.

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An Example
Determine the E?cell for the reaction 2 Al (s)
3 Cu2 (aq) ? 2 Al3 (aq) 3 Cu (s)
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Strengths Of Oxidizing And Reducing Agents

• For Table 18.1
• All of the reactions are written as reduction
  half-reactions.
• The more positive Eo, the stronger the oxidizing
  agent.
• The more negative Eo, the stronger the reducing
  agent.

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Criteria For Spontaneous Change In Redox
Reactions

• In general ?G -n F Ecell
• At standard conditions ?Go -n F Eocell
• If Ecell gt0, the reaction is spontaneous.
• If Ecell lt0, the reaction is non-spontaneous.
• If Ecell 0, the system is at equilibrium.
• When a cell reaction is reversed, Ecell and ?G
  change signs.

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Criteria For Spontaneous Change In Redox
Reactions An Example

• Exercise 18.6 A
• Should the following reaction occur
  spontaneously as written for standard-state
  conditions?
• Cu2 (aq) 2 Fe2 (aq) ? 2 Fe3 (aq) Cu (s)

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Equilibrium Constants For Redox Reactions

• ?Go -RTlnKeq ?Go - nFEocell
• RTlnKeq nFEocell
• RT
• or Eocell ______ lnKeq
• nF
• R - gas constant 8.3145 J mol-1 K-1
• T - Kelvin temperature
• n - the number of moles of electrons involved in
  the reaction
• F - the faraday constant 96,485 coulombs mole-1

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Equilibrium Constants For Redox Reactions An
Example

• Exercise 18.8A
• Calculate the values of ?Go and Kep at 25oC for
• 3 Mg (g) 2 Al3 (1M) ? 3 Mg2 (aq) 2 Al (s)

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Summarizing The Important Relationships
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Putting Them Together
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Corrosion Metal Loss Through Voltaic Cells

• In moist air, iron can be oxidized to Fe2,
  particularly at scratches, nicks, or dents. These
  areas are referred to anodic areas.
• Other regions of the iron serve as cathodic
  areas, where the electrons from the anodic areas
  reduce atmospheric oxygen to the OH- ion.
• Iron (II) ions migrate from the anodic areas to
  the cathodic areas where they combine with the
  hydroxide ions. The iron (II) is then further
  oxidized to iron (III) by atmospheric oxygen.
  Fe2O3xH2O is common rust.

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Corrosion Of An Iron Piling
Overall 2Fe(s) O2 (g) 2H2O (l) ? 2Fe2(aq)
4OH-(aq)
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Protecting Iron From Corrosion

• The simplest line of defense against the
  corrosion of iron is to paint it to exclude
  oxygen from the surface.
• Another approach is to coat the iron with a thin
  layer of a less active metal.
• An entirely different approach is to protect iron
  with a more active metal, as in the zinc-clad
  iron known as galvanized iron.
• One other approach to protecting iron, similar to
  galvanization, is cathodic protection. The iron
  object to be protected is connected to a chunk of
  an active metal and the iron serves as the
  reduction half-cell.

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Summary

• An oxidation-reduction reaction can be separated
  into two half-reactions, one for oxidation and
  one for reduction.
• Half-reactions can be conducted in half-cells.
• A cell diagram for a voltaic cell is written with
  the anode on the left and the cathode on the
  right.
• A standard hydrogen electrode has H ion at unit
  activity in equilibrium with H2 gas on an inert
  platinum electrode.
• Standard cell potential can be calculated from
  standard potentials of cathode and anode.
• Eocell Eo(cathode) Eo(anode)

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Summary (Continued)

• A redox reaction for which Ecell gt 0 occurs
  spontaneously.
• ?Go, Eocell and Kep are related
• ?Go -RTlnKeq - nFEocell
• A corroding metal consists of anodic areas, at
  which dissolution of the metal occurs, and
  cathodic areas, where atmospheric oxygen is
  reduced to hydroxide ion.
• A metal can be protected against corrosion by
  plating it with a second metal that corrodes less
  readily.