Electrochemistry Chap 17

1
Electrochemistry - Chap 17

• Galvanic Cells
• Standard Reduction Potentials
• Cell Potential, Electrical Work, and Free Energy
• Dependence of Cell potential on Concentration
• Batteries
• Corrosion
• Electrolysis
• Electrolytic Processes

2
Review of Terms

• oxidation-reduction (redox) reaction
  involves a transfer of electrons from the
  reducing agent to the oxidizing agent.
• oxidation loss of electrons
• reduction gain of electrons

3
Review - Steps in Balancing Oxidation-Reduction

• Equations in Acidic solutions
• 1. Assign oxidation numbers to each atom so that
  you know what is oxidized and what is reduced
• 2. Split the skeleton equation into two
  half-reactions-one for the oxidation reaction
  (element increases in oxidation number) and one
  for the reduction (element decreases in oxidation
  number)
• Fe2 (aq) ? Fe3 (aq)
• MnO4- (aq) ? Mn2 (aq)

4
Review - Steps in Balancing Oxidation-Reduction

• 3. Complete and balance each half reaction
• Fe2(aq) ? Fe3(aq)
• MnO4-(aq) ? Mn2(aq)
• a. Balance all atoms except O and H
• b. Balance O atoms by adding H2O to one side of
  the equation
• MnO4-(aq) ? Mn2(aq) 4 H2O
• c. Balance H atoms by adding H to one side of
  the equation
• 8 H(aq) MnO4-(aq) ? Mn2(aq) 4 H2O
• d. Balance the electric charge by adding
  electrons (e-) to the more positive side
• Fe2(aq) ? Fe3(aq) e- oxidation
• 8 H(aq) MnO4-(aq) 5 e- ? Mn2(aq) 4 H2O
  reduction

5
Review - Steps in Balancing Oxidation-Reduction

• 4. Combine the two half-reactions to obtain the
  balanced oxidation-reduction equation
• a. Multiply each half reaction by a factor which
  will allow the e-s to cancel when the equations
  are added
• 5 Fe2(aq) ? 5 Fe3(aq) 5 e- oxidation
• 8 H(aq) MnO4-(aq) 5 e- ? Mn2(aq) 4 H2O
  reduction
• b. Simplify the equation by canceling species
  which occur on both sides of the equation and
  reduce the coeffi-
• cients to the smallest whole number.
• Check
• 5 Fe2(aq) 8 H(aq) MnO4-(aq) ?
• 5 Fe3(aq) Mn2(aq) 4 H2O

6
Review - Additional Steps for Balancing Oxidation
Reduction Equations in Basic Solutions

• 5. Add the same number of OH- ions to both sides
  of the equation as H ions
• 6. Note that when H reacts with OH-, it forms
  H2O. Cancel any H2Os that occur on both sides of
  the equation and reduce the equation to simplest
  terms

7
Electrochemistry

• The study of the interchange of chemical and
  electrical energy.
• Which chemically is redox reactions!!!

8
Half-Reactions

• The overall reaction is split into two
  half-reactions, one involving oxidation and one
  reduction.
• 8H MnO4? 5Fe2 ? Mn2 5Fe3 4H2O
• Reduction 8H MnO4? 5e? ? Mn2 4H2O
• Oxidation 5Fe2 ? 5Fe3 5e?

9
Galvanic Cell

• A device in which chemical energy is changed to
  electrical energy.
• Spontaneous process
• Common use? Batteries!!!

10
Galvanic Cell

• a galvanic cell uses a spontaneous redox reaction
  to produce a current that can be used to do work
• The oxidation reaction occurs at the anode
  (negative electrode)
• The reduction reaction occurs at the cathode
  (positive electrode)

11
Galvanic Cell
12
Galvanic Cell
Typical reaction Anode M ? M e- Cathode
M e- ? M
13
Cell Potential

• Cell Potential or Electromotive Force (emf)
  The pull or driving force on the electrons.
• the unit of electrical potential is the volt (V)
  which is defined as 1 joule of work per coulomb
  of charge transferred

14
Standard Reduction Potentials

• the reactions in a galvanic cell are always
  oxidation-reduction reactions that can be written
  into two half-reactions
• We can then assign the potential for each of the
  half-reactions so that when the cell is
  constructed for a given pair of half reactions,
  we obtain the cell potential by summing the
  potentials
• the standard hydrogen electrodes potential is
  arbitrarily set to zero volts

15
Standard Reduction Potentials

• The E? values corresponding to reduction
  half-reactions with all solutes at 1M and all
  gases at 1 atm.
• Cu2 2e? ? Cu
• E? 0.34 V vs. SHE
• SO42? 4H 2e? ? H2SO3 H2O
• E? 0.20 V vs. SHE

16
Line notation

• A short code for an electrochemical cell
  Line Notation
• Cu2 2e? ? Cu
• E? 0.34 V vs. SHE
• SO42? 4H 2e? ? H2SO3 H2O
• E? 0.20 V vs. SHE
• PtH(aq), SO42?(aq), H2SO3 (aq)Cu2(aq), Cu(s)
• The anode components are listed to the left side
  of the double line and cathode are on the right
  side of the double line. The electrode is listed
  on either end of the equation separated by a
  single line if it is a different material (see
  the platnum electrode above)

17
emf and Work
From the above equations you can calculate the
maximum cell potential and work Any real
spontaneous process actual work realized is
less than the calculated maximum
18
Free Energy and Cell Potential

• ?G? ?nF??
• n number of moles of electrons
• F Faraday 96,485 coulombs per mole of
  electrons
• So we can in turn relate work to free energy for
  a galvanic cell

19
Standard Conditions

• When using standard conditions all concentrations
  are 1M
• What happens when we place a higher concentration
  of one of the components
• Le Châteliers principle will drive the
  reaction

20
Concentration Cell

• Because cell potentials are dependent on
  concentration
• a cell in which both compartments have the same
  components but at different concentrations can be
  made!!

21
The Nernst Equation

• We can calculate the potential of a cell in which
  some or all of the components are not in their
  standard states.

Free energy equation from chap 16
Using the above equation and substituting into
the first equation yields eq3
Doing some simple algebra yields eq4
Eq4 at 25oC yields eq5
22
Calculation of Equilibrium Constants for Redox
Reactions

• At equilibrium, Ecell 0 and Q K.

23
Batteries

• A battery is a galvanic cell or, more commonly, a
  group of galvanic cells connected in series.
• Look at the various examples in the book Lead,
  Dry cell and Fuel cells

24
Fuel Cells

• . . . galvanic cells for which the reactants are
  continuously supplied.
• 2H2(g) O2(g) ? 2H2O(l)
• anode 2H2 4OH? ? 4H2O 4e?
• cathode 4e? O2 2H2O ? 4OH?

25
Corrosion

• Can be viewed as returning metal to their natural
  state.
• Metals corrode because the oxidize easily
• Different metals corrode at different rates
• Due to anomalies among the metals
• Al forms Al2O3 a thin film that protects the
  rest of the metal
• Iron forms various oxides that flake off and
  expose a new layer of metal

26
Corrosion

• Some metals, such as copper, gold, silver and
  platinum, are relatively difficult to oxidize.
  These are often called noble metals.

27
Corrosion of Iron

• Iron or steel has a non-uniform surface
• Anodic surface and cathodic surface
• 4Fe2 O2 (42n)H2O ? 2Fe2O3.nH2O 8H

28
Electrolytic Cell

• in an electrolytic cell electrical energy is used
  to produce a chemical change
• the process of electrolysis involves forcing a
  current through a cell to produce a chemical
  change for which the cell potential is negative
• electrolysis has great practical importance
  examples, charging a battery, producing aluminum
  metal, and chrome plating

29
Electrolytic Plating

• Disposition of neutral metal on the electrode
  from ions in the solution of the electrolytic
  cell
• Process used in Chromium, Silver, Gold, and
  Copper plating

30
Stoichiometry of Electrolysis

• How much chemical change occurs with the flow of
  a given current for a specified time?
• current and time ? quantity of charge ?
• moles of electrons ? moles of analyte ?
• grams of analyte