1. Electrochemistry Redox Reactions and Electrochemical Cells
Review of Oxidation Reduction Concepts
Half-Reaction Method for Balancing Redox Reactions
Electrochemical Cells
Voltaic Cell: Using Spontaneous Reactions to Generate Electrical Energy
Construction and Operation
Cell Notation
Why Does the Cell Work
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2. Electrochemistry Cell Potential: Output of a Voltaic Cell
Standard Cell Potentials
Strengths of Oxidizing and Reducing Agents
Free Energy and Electrical Work
Standard Cell Potential
Effect of Concentration of Ecell
Changes in Ecell During Cell Operation
Concentration Cells
Electrochemical Processes in Batteries
Primary (Nonrechargeable Batteries)
Fuel Cells
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3. Electrochemistry Corrosion: A case of Environmental Electrochemistry
Corrosion of Iron
Protecting Against Corrosion
Electrolytic Cells: Using electrical energy to drive Nonspontaneous Reactions
Construction and Operation
Predicting Electrolysis Products
Stoichiometry of Electrolytes
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4. Electrochemistry Electrochemistry
The study of the relationship between chemical change (reactions) and the flow of electrons (electrical work)
Electrochemical Systems
Electrolytic – Work done by absorbing free energy from a source (passage of an electrical current through a solution) to drive a nonspontaneous reaction
Voltaic/Galvanic – Release of free energy from a spontaneous reaction to produce electricity (Batteries)
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5. Electrochemistry Oxidation-Reduction Concepts Review
Oxidation – Loss of Electrons
Reduction – Gain of Electrons
Oxidizing Agent – Species that causes another species to be oxidized (lose electrons)
Oxidizing agent is reduced (gains e-)
Reducing Agent – Species that cause another species to be reduced (gain electrons)
Reducing agent is oxidized (loses e-)
Oxidation (e- loss) always accompanies Reduction (e- gain)
Total number of electrons gained by the atoms/ions of the oxidizing agent always equals the total number of electrons lost by the reducing agent
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6. Electrochemistry
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7. Electrochemistry Oxidation Number
A number equal to the magnitude of the charge an atom would have if its shared electrons were held completely by the atom that attracts them more strongly
The oxidation number in a binary ionic compound equals the ionic charge
The oxidation number for each element in a covalent compound (or polyatomic ion) are assigned according to the relative attraction of an atom for electrons
See next slide for a summary of the rules for assigning oxidation numbers
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8. Electrochemistry
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9. Electrochemistry Balancing Redox Reactions Oxidation Number Method
Half-Reaction Method
The balancing process must insure that:
The number of electrons lost by the
reducing agent equals the number of
electrons gained by the oxidizing agent
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10. Electrochemistry Oxidation Number Method
Assign oxidation numbers to all elements in the reaction
From changes in oxidation number of given elements, identify oxidized and reduced species
For each element that undergoes a change of oxidation number, compute the number of electrons lost in the oxidation and gained in the reduction from the oxidation number change (Draw tie-lines between these atoms)
Multiply one or both these number by appropriate factors to make the electrons lost equal to the electrons gained
Use factors as coefficients in reaction equation
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11. Electrochemistry Half-Reaction Method
Applicable to Acid or Base solutions
Does not usually require Oxidation Numbers (ON)
Procedure
Divide the overall reaction into:
Oxidation Half-Reaction
Reduction Half-Reaction
Balance each half-reaction for atoms & charge
Multiply one or both reactions by some integer to make electrons gained equal to electrons lost
Recombine to given balanced redox equation
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12. Electrochemistry Redox Half-Reaction Method – Example
Divide steps into Half-Reactions
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13. Electrochemistry Balance Atoms & Charges for Cr2O72- / Cr3+
Balance Atoms & Charges for I- / I2
Add 7 Water molecules
to balance Oxygen
Add 14 H+ ions on left to balance 14 H on right
Add 6 electrons (e-) on left to balance reaction charges
(6 electrons gained  this is the reduction reaction
No need to add H2O or H+
Add 2 electrons (e-) on right to balance reaction charges
(2 electrons lost  this is the oxidation reaction
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14. Electrochemistry Redox Half-Reaction Method – Example (con’t)
Multiply each half-reaction, if necessary, by an integer to balance electrons lost/gained
2 e- lost in oxidation reaction and 6 e- gained in reduction
 Multiply oxidation half-reaction by 3
Add 2 half-reactions together
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15. Sodium Permanganate & Sodium Oxalate
Electrochemistry
Half-Reaction Method in a “Basic” solution
Sodium Permanganate & Sodium Oxalate
NaMnO Na2C2O4
Half-Reactions
Multiply each reaction by appropriate integer
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16. Electrochemistry Sodium Permanganate & Sodium Oxalate (con’t)
Add reactions
Add OH- to neutralize H+ , balance H2O, and form “basic” solution
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17. Electrochemistry Electrochemical Cells Voltaic (Galvanic) Cells
Use spontaneous reaction (G < 0) to generate electrical energy
Difference in Chemical Potential energy between higher energy reactants and lower energy products is converted to electrical energy to power electrical devices
Thermodynamically - The system does work on the surroundings
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18. Electrochemistry Electrochemical Cells Electrolytic Cells
Uses electrical energy to drive nonspontaneous reaction (G > 0)
Electrical energy from an external power supply converts lower energy reactants to higher energy products
Thermodynamically – The surroundings do work on the system
Examples – Electroplating and recovering metals from ores
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19. Electrochemistry
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20. Zn(s) Zn2+(aq) Cu2+(aq) Cu(s)
Electrochemistry
Electrochemical Cells
Cell notation is used to describe the structure of a voltaic (galvanic) cell
For the Zn/Cu cell, the cell notation is:
Zn(s) Zn2+(aq) Cu2+(aq) Cu(s)
= phase boundary (solid Zn vs. Aqueous Zn2+)
= salt bridge
Anode reaction (oxidation) is left of the salt bridge
Cathode reaction (reduction) is right of the salt bridge
Half-cell components usually appear in the same order as in the half-reactions (Zn(s) + 2e-  Zn2+).
Zinc solid loses 2 e- (oxidized) to produce zinc(II) at the negative ANODE
Copper(II) gains 2e- (reduced) to form copper metal at positive CATHODE
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21. Electrochemistry Voltaic (Galvanic) Cells
Zinc metal (Zn) in solution of Cu++ ions
Construction of a Voltaic Cell
The oxidizing agent (Zn) and reducing agent (Cu2+) in the same beaker will not generate electrical energy
Separate the half-reactions by a barrier and connect them via an external circuit (wire)
Set up salt bridge between chambers to maintain neutral charge in electrolyte solutions
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22. Electrochemistry Oxidation Half-Cell
Anode Compartment – Oxidation of Zinc (An Ox)
Zinc metal in solution of Zn2+ electrolyte (ZnSO4)
Zn is reactant in oxidation half-reaction
Conducts released electrons (e-) out of its half-cell
Reduction Half-Cell
Cathode Compartment – Reduction of Copper (Red Cat)
Copper bar in solution of Cu2+ electrolyte (CuSO4)
Copper metal is product in reduction half-cell reaction
Conducts electrons into its half-cell
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23. Electrochemistry
Zinc-Copper Voltaic Cell
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24. Electrochemistry Relative Charges on the Anode/Cathode electrodes
Electrode charges are determined by the source of the electrons and the direction of electron flow
Zinc atoms are oxidized (lose 2 e-) to form Zn2+ at the anode
Anode – negative charge (e- rich)
Released electrons flow to right toward cathode to be accepted by Cu2+ to form Cu(s)
Cathode – positive charge (e- deficient)
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25. Electrochemistry Purpose of Salt Bridge
Electrons from oxidation of Zn leave neutral ZnSO4 solution producing net positive charge
Incoming electrons to CuSO4 solution would produce net negative charge in solution as copper ions are reduced to copper metal
Resulting charge imbalance would stop reaction
Salt bridge provides “liquid wire” allowing ions to flow through both compartments completing circuit
Salt bridge constructed of an inverted “U-tube” containing a solution of non-reacting Na+ & SO42- ions in a gel
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26. Electrochemistry Active vs Inactive Electrodes Active Electrodes
Electrodes in Zn/Cu2+ cell are active
Zinc & Copper bars are components of the cell reactions
Mass of Zn bar decreases as Zn2+ ions in cell solution increase
Mass of Copper bar increases as Cu2+ ions accept electron to form more copper metal
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27. Electrochemistry Active vs Inactive Electrodes Inactive Electrodes
In many Redox reactions, one or the other reactant/product is not capable of serving as an electrode
Inactive electrodes - Graphite or Platinum
Can conduct electrons into and out of half-cells
Cannot take part in the half-reactions
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28. Inactive Graphite Electrodes
Electrochemistry
Voltaic Cell
with
Inactive Graphite Electrodes
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29. Electrochemistry Cell Potential
The movement of electrons is analogous to the pumping of water from one point to another
Water moves from a point of high pressure to a point of lower pressure. Thus, a pressure difference is required
The work expended in moving the water through a pipe depends on the volume of water and the pressure difference
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30. Electrochemistry Cell Potential Movement of Electrons
An electric charge moves from a point of high electrical potential (high electrical pressure) to one of lower electrical potential
The work expended in moving the electrical charge through a conductor depends on the potential difference and the amount of charge
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31. Electrochemistry Cell Potential
Purpose of a voltaic cell is to convert the free energy of a spontaneous reaction into the kinetic energy of electrons moving through an external circuit (electrical energy)
Electrical energy is proportional to the difference in the electrical potential between the two cell electrodes
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32. Electrochemistry Cell Potential
Positive Cell Potential – Electrons flow spontaneously from the negative electrode (Anode) to the positive electrode (Cathode)
Negative cell potential – is associated with a “nonspontaneous” cell reaction
Cell potential for a cell reaction at equilibrium would be “0”
As with Entropy, there is a clear relationship between Ecell , K, and G
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33. Electrochemistry Units of Cell Potential
The SI (metric) unit of electrical charge is the:
Coulomb (C)
The SI (metric) unit of current is the:
Ampere (A)
The SI (metric) unit of electrical potential is the:
“Volt (V)”
By definition, the energy released by a potential difference of one volt moving between the anode and cathode of a voltaic cell releases 1 joule of work per coulomb of charge
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34. Electrochemistry
The charge (F) that flows through a cell equals the number of moles of electrons (n) transferred times the charge of 1 mol of electrons
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35. Electrochemistry Standard Cell Potential
Eocell – The potential measured at a specific temperature (298 K) with no current flowing and all concentrations in their “Standard States”
1 atm for gases
1 M for solutions
Pure solids for electrodes
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36. Electrochemistry Standard Electrode Half-Cell Potentials
Eohalf-cell – Potential associated with a given half-cell reaction (electrode compartment) when all components are in “Standard States”
Standard Electrode Potential for a half-cell reaction, whether anode (oxidation) or cathode (reduction) is written as a “reduction”
Ex.
would be written:
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37. Electrochemistry Standard Electrode Half-Cell Potentials
Electrons flow spontaneously from Anode (negative) to Cathode (positive)
Cathode must have a more “Positive” Eohalf-cell than the Anode
For a “positive” Eocell
The standard cell potential is the difference between the standard electrode potential of the “Cathode” (reduction) half-cell and the standard electrode potential of the “Anode” (oxidation) half-cell
Standard half-cell potentials are “intensive” properties, thus their values do NOT have to be adjusted for stoichiometry (# of moles)
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38. Electrochemistry The Standard Hydrogen Electrode
Half-cell potentials are not absolute quantities
The values found in tables are determined relative to a “Standard”
The Standard Electrode potential is defined as zero
(Eoreference) = 0.00
The “standard reference half-cell” is a standard
“Hydrogen” electrode
Specially prepared Platinum electrode immersed in a 1 M aqueous solution of a strong acid through which H2 gas at 1 atm is bubbled
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39. Electrochemistry Reference Half-Cell and Unknown Half-Cell
The “Standard” electrode can act as either the “Anode” or the “Cathode”
Oxidation of H2 (lose e-) at anode half-cell and reduction of unknown at cathode half-cell
Reduction of H+ (gain e-) at cathode half-cell and oxidation of unknown at anode half-cell
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40. Electrochemistry Table of Standard Electrode Potentials
(The emf Series)
All Values are relative to the “standard hydrogen (reference) electrode
All reactions are written as “reductions”
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41. Electrochemistry EMF Series
All Values are relative to the “standard hydrogen (reference) electrode
All reactions are written as “reductions”
F2 is strongest oxidizing agent (high, positive Eo)
Fluorine is very electronegative
It is easily reduced (gain e-) to form weak reducing agent, F- (reluctant to lose electrons)
Li metal is strongest reducing agent (low, negative Eo)
has low ionization potential
easily oxidized (loses e-) to form strong oxidizing agent, Li+ (reluctant to gain electrons)
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42. Electrochemistry Writing Spontaneous Redox Reactions
Similarities – Acid/Base vs Redox
Acid Strength vs Base Strength using Ka & Kb values
Redox (Oxidizing agent vs Reducing agent) using Eo values
Tables of Standard Electrode Potentials (Eo)
The stronger oxidizing agent (species on left side of table) has a half-reaction with a larger (more positive or less negative) Eo
The stronger reducing agent (species on the right side of table) has a half-reaction with a smaller (less positive or more negative) value
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43. Electrochemistry Writing Spontaneous Redox Reactions
A spontaneous reaction (Eocell > 0) will occur between an oxidizing agent and any reducing agent that lies below it in the table
Zn(s) (reducing agent on right side of table with Eo = V) will react spontaneously with Cu2+ (oxidizing agent of left side of table with Eo = V) which lies above Zn in the table
A spontaneous reaction will occur when the half- reaction higher in the list occurs at the cathode (reduction) as written and the half-reaction lower in list occurs at the anode (oxidation) in reverse
Recall all half-reactions are written as “reductions” in the electrode potential table
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44. Electrochemistry Relative Reactivities of Metals
Metals that displace H2 from acid
If the Eocell for the reaction of H+ is more positive for metal A than it is for metal B, metal A is a stronger reducing agent than metal B and a more active metal
Metals Li through Pb (includes Fe) in the standard electrode potential list (appendix D) lie below H+ and give positive Eocell when reducing H+ to H2, i.e., Hydrogen gas is released
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45. Electrochemistry Relative Reactivities of Metals
Metals that cannot displace H2 from acid
Metals that lie above the standard hydrogen reference half-reaction cannot reduce H+ from acids
The Eocell for the reversed metal half-reaction is negative and the reaction does not occur
The higher the metal in the list, the more negative is its Eocell for the reduction of H+ to H2, thus its reducing strength (and reactivity) is less
Thus, Gold (Au3+, Eo = +1.5V) is less active than Silver (Ag+, Eo = +0.8V) and does not release Hydrogen gas
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46. Electrochemistry Relative Reactivities of Metals
Metals that displace H2 from water
Metals that lie below the half-cell reaction potential for water can displace H2 from water
In the reaction below the E value for water is not the standard state value listed in the table because in pure water, [OH-] is 1.0 x 10-7 M, not the standard state value of 1 M (-0.83 V)
Ecell > 0  Sodium displaces Hydrogen from water
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47. Electrochemistry Relative Reactivities of Metals
Metals that can displace other metals from solution
Any metal that is lower in the standard electrode half-cell list can reduce the ion of a metal that is higher in the list, thus displacing that metal from solution (See next slide and slide #48)
Ecell > 0 Zinc is the stronger reducing agent reducing Fe2+ to Fe and displacing it from solution
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48. Electrochemistry
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49. Electrochemistry Free Energy and Electrical Work Electrical Work
Potential (Ecell, in volts) times the charge
Ecell measured with no current flowing
No energy lost to heating
Ecell voltage is maximum possible for cell
Work is maximum possible
Only reversible process can do maximum work
Reversible process with no current flow:
Forward reaction if opposing potential is smaller
Reverse reaction if opposing potential is larger
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50. Electrochemistry Spontaneous Reaction – G < 0
Spontaneous Reaction – Ecell > 0
The voltaic cell loses energy as it does work on the surroundings; thus the work term (wmax) is negative
(Recall slide # 39)
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51. Electrochemistry Electrical Work (con’t)
Relate standard cell potential to equilibrium constant (K) of the redox reaction
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52. Electrochemistry
Summary Relationship between
Go Eocell K
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53. Electrochemistry Effect of Concentration on Cell Potential
Most cells do not start with concentrations in their “standard” states
Recall:
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54. Electrochemistry Changes in Potential During Cell Operation
The potential of a cell changes as the concentration of the cell components change
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55. Electrochemistry
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56. Electrochemistry Concentration Cells
In a cell composed of the same substance, but differing concentrations in the two half-cells, the two concentrations move to equilibrate producing electrical energy
The cell reaction is the “sum” of identical half-cell reactions written in opposite directions
The Standard Electrode Potentials (Eocell) are both based on a 1 M solution (standard conditions), so they “cancel” each other, i. e., Eocell = 0
The non-standard cell potential, Ecell, depends on the ratio of the two concentrations [A]dil / [A]conc = Q
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57. Ecell decreases until Ecell = 0 (Q = K)
Electrochemistry
How the Concentration Cell Works
The dilute solution is in the Anode compartment (oxidation) and the concentrated solution is in the Cathode compartment (Reduction)
In the Anode (dilute) half-cell, Cu atoms give up 2 electrons and the resulting Cu2+ ions enter the solution and make it more concentrated
In the Cathode (conc) half-cell, Cu2+ ions gain 2 electrons and the resulting Cu atoms plate out on the electrode, making the solution less concentrated
In this type of Voltaic cell, the dilution continues until equilibrium is attained, i.e.,
Ecell decreases until Ecell = 0 (Q = K)
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58. Electrochemistry Alkaline Battery Nickel-Metal Hydride (Ni-MH) Battery
Lithium Ion Battery
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