1. Electrochemistry MAE-212
Dr. Marc Madou, UCI, Winter 2018
Class I Thermodynamics of the Electromotive Force (I)
Constructing a Daniell Cell

2. Table of Content
Electrochemistry Definitions Electrochemical Cell Terminology Standard Electrode Potentials Ecell, ΔG, and Keq
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3. Electrochemistry Definitions
Electron transfer reactions are oxidation-reduction or redox reactions.
They can either result in the generation of an electric current (electricity) or they can be caused by imposing an electric current.
Therefore, this field of chemistry is often called ELECTROCHEMISTRY. In other words in electrochemical reactions, electrons are transferred from one species to another.
In order to keep track of what loses electrons and what gains them, we assign oxidation numbers.
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4. Electrochemistry Definitions
A species is oxidized when it loses electrons.
Here, zinc loses two electrons to go from neutral zinc metal to the Zn2+ ion.
Below we show two other types of half-cells:
Cu(s) + Zn2+(aq)
No reaction
Cu(s) + 2Ag+(aq)
Cu2+(aq) + 2 Ag(s)
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5. Electrochemistry Definitions
A species is reduced when it gains electrons.
Here, each of the H+ ‘s gains an electron, and they combine to form H2.
In a half-cell there is only one electrode: both reaction occur on the same surface
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6. Electrochemistry Definitions
What is reduced is the oxidizing agent.
H+ oxidizes Zn by taking electrons from it.
What is oxidized is the reducing agent.
Zn reduces H+ by giving it electrons.
In an electrochemical cell there are two electrodes in separate (ideally) half- cell:
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7. Electrochemistry Definitions
Each side in an electrochemical cell is a half-cell
Electrons flow from oxidation side to reduction side – determine which is which
Assignment of sign is as follows:
Negative terminal = oxidation (anode)
Positive terminal = reduction (cathode)
A salt bridge allows ions to move between cells so that a charge build up does not occur. This completes the circuit.
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8. Electrochemistry Definitions
Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons. Reduction has to occur at the cost of oxidation
LEO the lion says GER!
ose
lectrons
xidation
ain
lectrons
eduction
GER!
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9. Electrochemistry Definitions
OIL RIG s s
xidation
ose
eduction
ain
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10. Electrochemical Cell Terminology
Electromotive force, Ecell.
The cell voltage or cell potential.
Cell diagram below,
Shows the components of the cell in a symbolic way.
Anode (where oxidation occurs) on the left.
Cathode (where reduction occurs) on the right.
Boundary between phases shown by |.
Boundary between half cells (usually a salt bridge) shown by ||.
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) Ecell = V
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11. Electrochemical Cell Terminology
Galvanic cells.
Produce electricity as a result of spontaneous reactions.
Electrolytic cells.
Non-spontaneous chemical change driven by electricity.
A redox couple, M|Mn+
A pair of species related by a change in number of e-.
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12. Standard Electrode Potentials, E°
Cell voltages, the potential differences between electrodes, are among the most precise scientific measurements.
The potential of an individual electrode is difficult to establish.
Therefore an arbitrary zero is chosen:
The Standard Hydrogen Electrode (SHE)
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13. Standard Electrode Potentials, E°
Chemistry 140 Fall 2002
Standard Electrode Potentials, E°
2 H+(a = 1) + 2 e-  H2(g, 1 bar) E° = 0 V
Pt|H2(g, 1 atm)|H+(a = 1)
Potential values are referenced to a standard hydrogen electrode (SHE).
By definition, the reduction potential for hydrogen is 0 V:
2 H+ (aq, 1M) + 2 e−  H2 (g, 1 atm)
The two vertical lines indicate three phases are present.
For simplicity we usually assume that a = 1 at [H+] = 1 M and replace 1 bar by 1 atm.
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14. Standard Electrode Potential, E°
E° defined by international agreement.
The tendency for a reduction process to occur at an electrode.
All ionic species present at a=1 (approximately 1 M).
All gases are at 1 bar (approximately 1 atm).
Where no metallic substance is indicated, the potential is established on an inert metallic electrode (ex. Pt).
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15. Standard Electrode Potential, E°
Cu2+(1M) + 2 e- → Cu(s) E°Cu2+/Cu = ?
Pt|H2(g, 1 bar)|H+(a = 1) || Cu2+(1 M)|Cu(s) E°cell = V
anode
cathode
Standard cell potential: the potential difference of a cell formed from two standard electrodes.
E°cell = E°cathode - E°anode
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16. Standard Electrode Potential, E°
The strongest oxidizers have the most positive reduction potentials.
The strongest reducers have the most negative reduction potentials.
The greater the difference between the two, the greater the voltage of the cell.
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17. Standard Electrode Potential, E°
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18. Ecell, ΔG, and Keq ΔG = -nFE ΔG° = -nFE° Cells do electrical work.
Moving electric charge.
The voltage of the whole cell is the electrical energy that it gives off, measured in volts (V)
The current is the rate at which electrons pass through the cell, measured in amperes (A)
Faraday constant, (with Q, charge of a single electron and NA , Avogadro’s number), F=QNA or F = 96,485 C mol-1
ΔG = -nFE
ΔG° = -nFE°
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19. Ecell, ΔG, and Keq ΔG = -nFE
n = number of electrons transferred
F = Faraday constant = 96,500 C/mol or 96,500 J/V-mol
Why negative? Spontaneous reactions have +E and – ΔG.
Volts cancel, units for ΔG are J/mol
Standard conditions: ΔG° = -nFE°
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20. Ecell, ΔG, and Keq
Nonstandard conditions – during the life of an electrochemical cell this is most common: to calculate potentials in this condition one needs the Nernst Equation
E = E ° - (RT/nF) lnQ, with Q the reaction quotient. A reaction quotient: Q is a function of the activities or concentrations of the chemical species involved in a chemical reaction. In the special case that the reaction is at equilibrium the reaction quotient is equal to the equilibrium constant.
For the general reaction: aA + bB < = > cC + dD the reaction quotient is expressed as: Q = (aC)c(aD)d/(aA)a(aB)b
Consider Zn(s) + Cu2+ → Zn Cu(s)
What is Q in this case?
What is E when the ions are both 1M?
What happens as Cu2+ decreases?
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21. Ecell, ΔG, and Keq
Same electrodes and solutions, different molarities.
How will this generate a voltage? Look at Nernst Equation. E = E ° - (RT/nF)lnQ
When will it stop?
Such a concentration cell is the basis for a pH meter and the regulation of heartbeat in mammals
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22. Ecell, ΔG, and Keq
When an electrochemical cell continues to discharge, E eventually reaches 0. At this point, because ΔG = -nFE, it follows that ΔG = 0.
Equilibrium!
Therefore, Q = Keq
Derivation see next class
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23. Ecell, ΔG, and Keq ΔG < 0 for spontaneous change.
Therefore E°cell > 0 because ΔGcell = -nFE°cell
E°cell > 0
Reaction proceeds spontaneously as written.
E°cell = 0
Reaction is at equilibrium.
E°cell < 0
Reaction proceeds in the reverse direction spontaneously.
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