Electrochemistry

1
Electrochemistry

• Chapter 20
• Brown-LeMay

2
Review of Redox Reactions

• Oxidation - refers to the loss of electrons by a
  molecule, atom or ion - LEO goes
• Reduction - refers to the gain of electrons by an
  molecule, atom or ion GER
• Chemical reactions in which the oxidation state
  of one or more substances changes are called
  oxidation-reduction reactions (or redox
  reactions)

3
Zn(s) 2H(aq)? Zn2(aq) H2(g)

• Zn 0, Zn2 2 (LEO) reducing agent
• H 1, H2 0 (GER) oxidizing agent
• Thus, the oxidation number of both the Zn(s) and
  H(aq) change during the course of the reaction,
  and so, this must be a redox reaction
• Review balancing redox-equations

4
Balancing Redox by Half Reactions

• Half reactions are a convenient way of separating
  oxidation and reduction reactions.
• Balance the titration of acidic solution of
  Na2C2O4 (colorless) with KMnO4(deep purple)
• 1st Write the incomplete ½ reactions
• MnO4(aq) ? Mn2(aq) is reduced (pale pink)
• C2O4(aq) ? CO2(g) is oxidized

5

• MnO4-(aq) ? Mn2(aq) 4H2O
• C2O42-(aq) ? 2CO2(g)
• Then bal H by adding H
• 8H MnO4-(aq) ? Mn2(aq) 4H2O
• C2O42-(aq) ? 2CO2(g)
• finish by balancing the es
• For the permanganate 7 left and 2 on the right
• 5e- 8H MnO4-(aq) ? Mn2(aq) 4H2O
• On the oxalate 2- on the right and o on the left
• C2O42-(aq) ? 2CO2(g) 2e-

6

• 2(5e- 8H MnO4-(aq) ? Mn2(aq) 4H2O)
• 5(C2O42-(aq) ? 2CO2(g) 2e-)
• 10e- 16H 2MnO4- (aq) ? 2Mn2(aq) 8H2O
• 5C2O42-(aq) ? 10CO2(g) 10e-
• 16H2 MnO4- 5C2O42- ? 2Mn2(aq)8H2O
• 
  10CO2
  

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Balancing Eq in Basic Solution

• The same method is used but OH- is added to
  neutralize the H used.
• The equation must again be simplified by
  canceling the terms on both sides of the equation.

8
Voltaic Cells

• Spontaneous redox reactions may be use to perform
  electrical work
• Voltaic or galvanic cells are devices that
  electron transfer occurs in an external circuit.

9
Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)

• Zn0 is spontaneously oxidized to Zn2
• Cu2 is spontaneously reduced to Cu0
• Oxidization half reaction at the anode
• Zn(s) ? Zn2(aq) 2e-
• Reduction half reaction at the cathode
• Cu2(aq) 2e- ? Cu(s)

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• As oxidization occurs, Zn is converted to Zn2
  and 2e-.
• The electrons flow toward the cathode, where they
  are used in the reduction reaction
• We expect the Zn electrode to lose mass
• Electrons flow from the anode to cathode so the
  anode is negative and the cathode is positive

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• Electrons cannot flow through the solution they
  have to be transported though an external wire.
• Anions and cations move through a porous barrier
  or salt bridge
• Cations move into the cathodic compartment to
  neutralize the excess negatively charged ions
  (cathode Cu2 2e- ? Cu, so the counter ion of
  Cu is in excess)

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• A build up of excess charge is avoided by
  movement of cations and anions through the salt
  bridge

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• The anions move into the anodic compartment to
  neutralize the excess Zn2 ions formed by the
  oxidation.
• Molecular View
• - Rules of voltaic cells
• at the anode electrons are products
• oxidization occurs
• at the cathode electrons are reactants
• reduction occurs

14
Cell EMF the driving force

• Reactions are spontaneous because the cathode has
  a lower electrical potential energy than the
  anode
• Potential difference difference in electrical
  potential measured in volts
• One volt is the potential difference required to
  impart one joule (J) of energy to a charge of one
  coulomb (C)
• 1V 1 J/C

15
Cell EMF the driving force

• Electromotive force (emf) is the force required
  to push electrons though the external circuit
• Cell potential Ecell is the emf of a cell
• Ecell gt 0 for a spontaneous reaction
• For 1 molar, 1 atm for gases at 250C(standard
  conditions), the standard emf
• (standard cell potential) E0cell

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Standard Reduction Potentials Cal. Cell
Potentials

• Standard Reduction Potentials E0red are measured
  relative to a standard
• The emf of a cell is
• E0cell E0red(cathode) E0red(anode)
• The standard hydrogen electrode is used as the
  standard (standard hydrogen electrode)
• (SHE)
• 2H(aq,1M)2e- ? H2(g,1atm) E0cell 0 V

17
The standard hydrogen electrode

• The (SHE) is assigned a potential of zero
• Consider the following half reaction
• Zn(s) ? Zn2(aq) 2e-
• We can measure E0cell relative to the SHE
• In the cell the SHE is the cathode
• It cons of a Pt electrode in a tube 1M Hsol
• H2 is bubbled through the tube
• E0cell E0red(cathode)-E0red(anode)
• 0.76V 0V-E0red(anode)
• Therefore E0red(anode) -0.76V

18
The standard hydrogen electrode

• Standard electrode potentials are written as
  reduction reactions
• Zn2(aq) (aq,1 M) 2e- ? Zn(s) E0 -0.76V
• Since the reduction potential is negative in the
  presence of the SHE the reduction of Zn2 is
  non-spontaneous
• However the oxidization of Zn2 is spontaneous
  with the SHE

19
Standard reduction potential

• The standard reduction potential is an intensive
  property
• Therefore, changing the stoichiometric
  coefficient does not affect E0red
• 2Zn2(aq) 4e- ? 2Zn(s) E0red -0.76 V

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• E0red gt 0 are spontaneous relative to the SHE
• E0red lt 0 are non- spontaneous relative to the
  SHE
• The larger the difference between E0red values
  the larger the E0cell
• The more positive the E0cell value the greater
  the driving force for reduction

21
Oxidizing and Reducing Agents

• Consider the table of standard reduction
  potentials
• We use the table to determine the relative
  strength of reducing and oxidizing agents
• The more positive the E0red the stronger the
  oxidizing agent (written as the reactant)
• The more negative the E0red the stronger the
  reducing agent (written as the product)

22
Oxidizing and Reducing Agents

• We can use tables to predict if one reactant can
  spontaneously oxidize or reduce another
• Example F2 can oxidize H2 or Li
• Ni2 can oxidize Al(s)
• Li can reduce F2

23
Spontaneity of Redox Reactions

• E0cell E0red(red process) E0red(oxid process)
• Consider the reaction
• Ni(s) 2Ag(aq) ? Ni2(aq) 2Ag(s)
• The standard cell potential is
• E0cell E0red (Ag/Ag) E0red(Ni2/Ni)
• E0cell (0.80 V) - (-0.28)
• E0cell 1.08 V the value indicates the reaction
  is spontaneous

24
EMF and free energy change

• Delta G -nFE
• where delta G is the change in free energy
• n the number of moles of electrons transferred
• F Faradays constant
• E emf of the cell

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EMF and free energy change

• F 96,500 C/mole- 96,500 J/(V)(mole-)
• Since n and F are positive, if Delta G lt 0, then
  E gt 0 and the reaction will be spontaneous.
• Effect of concentration on cell EMF
• the cell is function until E0 at which
  point equilibrium has been reached and the cell
  is dead
• The point at which E0 is determined by the
  concentrations of the species involved in the
  redox reaction

26
Walter Nernst (Nobel Prize 1920)

• Nernst Equation
• G G0 RTlnQ
• -nFE -nFE0 RTlnQ
• Solve the equation for
• E give the Nernst Eq
• E E0 - RT/nF lnQ
• Or for base 10 log
• E E0- 2.3RT/nF logQ

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• The nernst eq at 298K
• E E0 0.0592/n log Q
• Consider if you may
• Zn(s) Cu2(aq) ? Zn2(aq) Cu(s)
• If Cu2 5.0M and Zn2 0.05 M
• Ecell 1.10 V 0.0592/2 log 0.05/5 1.16 V
• Cell emf and chemical equilibrium
• Log K nE0/0.0592
• thus if we know cell emf, we can calc the
  equilibrium constant