Transcript Electrochemistry MAE-212 Dr. Marc Madou, UCI, Winter 2015 Class I Thermodynamics of the Electromotive Force (I) Constructing a Daniell Cell.

Electrochemistry MAE-212
Dr. Marc Madou, UCI, Winter 2015
Class I Thermodynamics of the Electromotive
Force (I)
Constructing a Daniell Cell
Table of Content
Electrochemistry Definitions
Electrochemical Cell Terminology
Standard Electrode Potentials
Ecell, ΔG, and Keq
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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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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) + 2Ag+(aq)
Cu2+(aq) + 2 Ag(s)
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Cu(s) + Zn2+(aq)
No reaction
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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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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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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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Electrochemistry Definitions
 Reduction (gaining electrons) can’t happen without an
oxidation to provide the electrons. Reduction has to
occur at the cost of oxidation
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Electrochemistry Definitions
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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)
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Ecell = 1.103 V
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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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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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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)
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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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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 = 0.340 V
anode
cathode
Standard cell potential: the potential difference of a cell formed from two standard
electrodes.
E°cell = E°cathode - E°anode
Electrode
1.Cathode:
2.Anode:
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Sign
+

Description
cations migrate to it.
cations reduced
anions migrate to it.
cations made.
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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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Standard Electrode Potential, E°
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Ecell, ΔG, and Keq
 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
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ΔG° = -nFE°
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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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At Equilibrium
Spontaneous
Not spontaneous
G
0

+
E
0
+

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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+ → Zn2+ + 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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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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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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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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