55
Lecture #2 of 18
56
Q: What’s in this set of lectures?
A: Introduction, Review, and B&F Chapter 1, 15 & 4 main concepts:
● Section 1.1: Redox reactions
● Chapter 15: Electrochemical instrumentation
● Section 1.2: Charging interfaces
● Section 1.3: Overview of electrochemical experiments
● Section 1.4: Mass transfer and Semi-empirical treatment of
electrochemical observations
● Chapter 4: Mass transfer
57
Looking forward… our review of Chapter “0”
● Cool applications
● Redox half-reactions
● Balancing electrochemical equations
● History of electrochemistry
● IUPAC terminology and Ecell = Ered – Eox
● Nernst equation and Common reference electrodes
● Standard and Absolute potentials
● Latimer and Pourbaix diagrams
● Calculating Ecell under non-standard state conditions
● Conventions
A Short History Lesson…
58
Electrochemistry associated with Luigi Galvani who discovered
“animal electricity,” while trying to Frankenstein frogs legs (1791)
Physician, Physicist, Philosopher
Luigi Galvani
(1737–1798)
from Wiki
59
Voltaic pile
Invented by Alessandro Volta (1800) but the elements of the pile
(galvanic cells) were named after Galvani.
Physicist
What are the half-reactions?
(salt water)
Alessandro Volta
(1745–1827)
from Wiki
Volta presenting his "Voltaic
Pile" to Napoleon and his
court… and now he is a Count!
At the Tempio Voltiano (the Volta Temple)
near Volta's home in Como, Italy.
http://en.wikipedia.org/wiki/Voltaic_pile
http://en.wikipedia.org/wiki/Alessandro_Volta
Galvanic Cells
60
Every non-equilibrium
cell is a galvanic cell (in
one direction, i.e. the
spontaneous direction)
Physically separating the
half-reactions allows the
electrons to go over a
long distance, from the
anode to the cathode via
a conductor: basis for
conversion of chemical
energy into electricity!
Salt bridge is an ionic conduit to
prevent buildup of charge in both
compartments and also to prevent
bulk mixing of the two solutions
Electrolysis of water
Chemist
Surgeon
William Nicholson
(1753–1815)
Sir Anthony Carlisle
(1768–1840)
Chemist
Chemist
Johann Wilhelm Ritter
(1776–1810)
William Cruickshank
(17??–1810(1))
61
Volta’s results were shared with
the scientific community and then,
boom, many people demonstrated
electrolysis the same year, and
later, electroplating!
http://en.wikipedia.org/wiki/Johann_Wilhelm_Ritter
62
Daniell (galvanic) Cell (1836)
anode
oxidation
No more H2 from the
(primary) battery!
cathode
reduction
Chemist, Meteorologist
Half-reactions are
physically separated!
Zn (s)  Zn2+ (aq) + 2e–
NET REACTION: Zn (s) +
Cu2+ (aq) + 2e–  Cu (s) John Frederic Daniell
Cu2+
(aq) 
(1790–1845)
Zn2+
(aq) + Cu (s)
from Wiki
Voltage Produced by Galvanic Cells
63
The difference in electric
potential between the anode
and the cathode is called:
 Cell potential
 Cell voltage
 emf (electromotive force)
Cell Diagram
Zn (s) + Cu2+ (aq)
This is horrible
It should be –0.76 V!
Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M and [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode
salt bridge cathode
64
EXAMPLE: What is the standard potential of an electrochemical
cell made of a Cd electrode in a 1.0 M Cd(NO3)2 solution and a
Cr electrode in a 1.0 M Cr(NO3)3 solution?
Which half-reaction is reducing?
Cd2+ (aq) + 2e–
Cr3+ (aq) + 3e–
Cd (s) E0 = -0.40 V Cd2+ will get reduced to Cd
Cr (s)
E0 = -0.74 V Cr will get oxidized to Cr3+
… thus, it is reducing
More negative of the two
Anode (oxidation):
Cr (s)
Cr3+ (1 M) + 3e– x 2
Cathode (reduction): 2e– + Cd2+ (1 M)
2Cr (s) + 3Cd2+ (1 M)
Cd (s)
x3
3Cd (s) + 2Cr3+ (1 M)
0
0 = E0
Ecell
–
E
cathode
anode
0 = -0.40 V – (-0.74 V)
Ecell
0 = +0.34 V (positive = spontaneous)
Ecell
The Daniell Cell (1836)
Chemist, Meteorologist
low impedance to measure current
–
–
Icell
John Frederic Daniell
(1790–1845)
from Wiki
Zn
Zn(s)
Cu
Zn2+(aq)
Cu2+(aq)
Zn | Zn2+(aq) || Cu2+(aq) | Cu
Cu(s)
The Daniell Cell (1836)
low impedance to measure current
–
–
Icell
This only works for
less than one second
and then stops due to
Kirchhoff’s current law,
which states that all
current at each location
must sum to zero…
… capacitive charging
Zn
Zn(s)
Cu
Zn2+(aq)
Cu2+(aq)
Zn | Zn2+(aq) || Cu2+(aq) | Cu
Cu(s)
The Daniell Cell (1836)
low impedance to measure current
–
–
Icell
Now it
works!
Ai+
Salt bridge
contains an
inert, redox
inactive salt
solution
(electrolyte)
Bi–
Zn
Zn(s)
Cu
Zn2+(aq)
Cu2+(aq)
Zn | Zn2+(aq) || Cu2+(aq) | Cu
Cu(s)
The Daniell Cell (1836)
low impedance to measure current
–
–
Icell
* Name this cell type
* Identify anode
* Identify cathode
* Name the
electrode signs
Ai+
Bi–
Zn
Zn(s)
Cu
Zn2+(aq)
Cu2+(aq)
Zn | Zn2+(aq) || Cu2+(aq) | Cu
Cu(s)
The Daniell Cell (1836)
low impedance to measure current
–
–
Icell
Ai+
* Name this cell type
* Identify anode
* Identify cathode
* Name the
electrode signs
* primary galvanic cell
Bi–
(A/–)
Zn(s)
Zn
Cu
Zn2+(aq)
Cu2+(aq)
Zn | Zn2+(aq) || Cu2+(aq) | Cu
(C/+)
Cu(s)
The Daniell Cell (1836)
This will eventually fully
discharge and reach
equilibrium (ΔG = Ecell = 0)
low impedance to measure current
–
–
Icell
Ai+
Then, either direction of
polarization bias results
in electrolytic function
(i.e. charging)!
Bi–
(A/–)
Zn(s)
Zn
Cu
Zn2+(aq)
Cu2+(aq)
Zn | Zn2+(aq) || Cu2+(aq) | Cu
(C/+)
Cu(s)
Electrochemistry:
conventions… oh, conventions!
Cathode – electrode where catholyte species are reduced
Anode – electrode where anolyte species are oxidized
71
Electrochemistry:
conventions… oh, conventions!
Cathode – electrode where catholyte species are reduced
Anode – electrode where anolyte species are oxidized
Negative/Positive electrode – cathode or anode?… it depends!
For the discharging (galvanic) battery, label the anode and cathode.
http://autoshop101.com/
72
Electrochemistry:
conventions… oh, conventions!
73
Cathode – electrode where catholyte species are reduced
Anode – electrode where anolyte species are oxidized
Negative/Positive electrode – cathode or anode?… it depends!
For the charging (electrolytic) battery, label the anode and cathode.
e–
e–
This is a generator
(AKA power supply)
cathode
anode
http://autoshop101.com/
Electrochemistry:
conventions… oh, conventions!
74
Cathode – electrode where catholyte species are reduced
Anode – electrode where anolyte species are oxidized
… Sheesh!
… Take-home message: For batteries, don’t call electrodes
anodes and cathodes (but convention used by most is for
discharging)
e–
e–
e–
anode
cathode
e–
anode
http://autoshop101.com/
cathode
PROBLEM TIME!
You try!
(a) What is the standard Ecell for a
galvanic cell based on zinc and silver?
(b) If we wanted to electrolytically
charge the cell from part a (before
any reactions took place), what
potential would we have to apply?
75
PROBLEM TIME!
You try!
(a) What is the standard Ecell for a
galvanic cell based on zinc and silver?
(b) If we wanted to electrolytically
charge the cell from part a (before
any reactions took place), what
potential would we have to apply?
(a) Ecell = +0.80 V – (-0.76 V) = 1.56 V
(b) Ebias < –1.56 V
76
International Union of Pure and Applied
Chemistry (IUPAC)
77
(Accepted) Nomenclature and Terminology that you know, but may have forgotten
•
Coulomb (in units of C = A·s) is the unit of charge (96,485 C are in a mole of singly
charged species = Faraday constant, F ≈ 96,500 C/mol ≈ 105 C/mol)
•
Electricity is the flow of current (A = C/s) and is negative (cathodic) or positive (anodic)
depending on the direction and sign of the current-carrying species (e.g. e–, H+)
•
(Electrode) (electric) potential (V or E; in units of V = J/C) is written as a reduction
This relates to Gibbs free energy as ΔG = –RT ln K = –nFEcell (electrical work per mole), and…
… partial molar Gibbs free energy is the electrochemical potential (𝜇, in units of J/mol)

Chemical potential (𝜇, in units of J/mol)

Galvani/Inner (electric) potential (ϕ, in units of V)
… and in summary, 𝜇 = 𝜇 + 𝑛𝐹ϕ
Also, standard state is a solvent, a solid, and a species at unit activity (~1 M solutes, ~1 bar gases)
Also, E = IR (Ohm’s law) when resistance is constant
•
Galvanic cells produce power (in units of W = A x V = C/s x J/C = J/s) by spontaneous
redox reactions
•
Electrolytic cells require a power input to drive redox reactions; thus, the reactions are
thermodynamically unfavorable
•
Batteries have anodes and cathodes, but these names change depending on if the
battery is being discharged (galvanic) or charged
(electrolytic),
and so
the terminology
(Sources:
B&F, M3LC
course
textbook,
of negative electrode and positive electrode is preferred
and http://goldbook.iupac.org/)
International Union of Pure and Applied
Chemistry (IUPAC)
78
(Accepted) Nomenclature and Terminology that you know, but may have forgotten
•
Coulomb (in units of C = A·s) is the unit of charge (96,485 C are in a mole of singly
charged species = Faraday constant, F ≈ 96,500 C/mol ≈ 105 C/mol)
integrate, over time
differentiate, with respect to time
•
Electricity is the flow of current (A = C/s) and is negative (cathodic) or positive (anodic)
depending on the direction and sign of the current-carrying species (e.g. e–, H+)
•
(Electrode) (electric) potential (V or E; in units of V = J/C) is written as a reduction
This relates to Gibbs free energy as ΔG = –RT ln K = –nFEcell (electrical work per mole), and…
… partial molar Gibbs free energy is the electrochemical potential (𝜇, in units of J/mol)

Chemical potential (𝜇, in units of J/mol)

Galvani/Inner (electric) potential (ϕ, in units of V)
… and in summary, 𝜇 = 𝜇 + 𝑛𝐹ϕ
Also, standard state is a solvent, a solid, and a species at unit activity (~1 M solutes, ~1 bar gases)
Also, E = IR (Ohm’s law) when resistance is constant
•
Galvanic cells produce power (in units of W = A x V = C/s x J/C = J/s) by spontaneous
redox reactions
•
Electrolytic cells require a power input to drive redox reactions; thus, the reactions are
thermodynamically unfavorable
•
Batteries have anodes and cathodes, but these names change depending on if the
battery is being discharged (galvanic) or charged
(electrolytic),
and so
the terminology
(Sources:
B&F, M3LC
course
textbook,
of negative electrode and positive electrode is preferred
and http://goldbook.iupac.org/)
International Union of Pure and Applied
Chemistry (IUPAC)
79
(Accepted) Nomenclature and Terminology that you know, but may have forgotten
•
Coulomb (in units of C = A·s) is the unit of charge (96,485 C are in a mole of singly
charged species = Faraday constant, F ≈ 96,500 C/mol ≈ 105 C/mol)
•
Electricity is the flow of current (A = C/s) and is negative (cathodic) or positive (anodic)
depending on the direction and sign of the current-carrying species (e.g. e–, H+)
•
(Electrode) (electric) potential (V or E; in units of V = J/C) is written as a reduction
This relates to Gibbs free energy as ΔG = –RT ln K = –nFEcell (electrical work per mole), and…
… partial molar Gibbs free energy is the electrochemical potential (𝜇, in units of J/mol)

Chemical potential (𝜇, in units of J/mol)
E0(Cu2+/0) = +0.34 V vs. SHE

Galvani/Inner (electric) potential (ϕ, in units of V)
E0(Cu0/2+) = –0.34 V vs. SHE
… and in summary, 𝜇 = 𝜇 + 𝑛𝐹ϕ
Also, standard state is a solvent, a solid, and a species at unit activity (~1
solutes,
~1 bar gases)
…M is
incorrect!
Also, E = IR (Ohm’s law) when resistance is constant
•
•
You can subtract redox
Galvanic cells produce power (in units of W = Apotentials
x V = C/s x but
J/C =
bychange
spontaneous
doJ/s)
not
the
redox reactions
sign of the potential and then
call itreactions;
an oxidation
Electrolytic cells require a power input to drive redox
thus, potential!
the reactions are
thermodynamically unfavorable
•
Batteries have anodes and cathodes, but these names change depending on if the
battery is being discharged (galvanic) or charged
(electrolytic),
and so
the terminology
(Sources:
B&F, M3LC
course
textbook,
of negative electrode and positive electrode is preferred
and http://goldbook.iupac.org/)
International Union of Pure and Applied
Chemistry (IUPAC)
80
(Accepted) Nomenclature and Terminology that you know, but may have forgotten
•
Coulomb (in units of C = A·s) is the unit of charge (96,485 C are in a mole of singly
charged species = Faraday constant, F ≈ 96,500 C/mol ≈ 105 C/mol)
•
Electricity is the flow of current (A = C/s) and is negative (cathodic) or positive (anodic)
depending on the direction and sign of the current-carrying species (e.g. e–, H+)
•
(Electrode) (electric) potential (V or E; in units of V = J/C) is written as a reduction
This relates to Gibbs free energy as ΔG = –RT ln K = –nFEcell (electrical work per mole), and…
… partial molar Gibbs free energy is the electrochemical potential (𝜇, in units of J/mol)

Chemical potential (𝜇, in units of J/mol)
E0(Cu2+/0) = +0.34 V vs. SHE

Galvani/Inner (electric) potential (ϕ, in units of V)
E0(Cu0/2+) = –0.34 V vs. SHE
… and
in summary,
𝜇 = 𝜇 + 𝑛𝐹ϕ it is best
Based on our
current
sign convention,
Also, standard
statepotentials;
is a solvent, ahowever,
solid, and aifspecies at unit activity (~1
solutes,
~1 bar gases)
…M is
incorrect!
to only write
reduction
Also, E = IR (Ohm’s law) when resistance is constant
we lived in an oxidation-potential-centric world,
You can subtract redox
we• could
write cells
themproduce
all (i.e.power
everything)
asW = Apotentials
Galvanic
(in units of
x V = C/s x but
J/C =
bychange
spontaneous
doJ/s)
not
the
redox
reactions
oxidation potentials; simply put, it is best to
sign of the potential and then
not• mixElectrolytic
the conventions
and
so
stick
with
call itreactions;
an oxidation
cells require a power input to drive redox
thus, potential!
the reactions are
reduction
potentials
thermodynamically unfavorable
•
Batteries have anodes and cathodes, but these names change depending on if the
battery is being discharged (galvanic) or charged
(electrolytic),
and so
the terminology
(Sources:
B&F, M3LC
course
textbook,
of negative electrode and positive electrode is preferred
and http://goldbook.iupac.org/)
International Union of Pure and Applied
Chemistry (IUPAC)
81
(Accepted) Nomenclature and Terminology that you know, but may have forgotten
•
Coulomb (in units of C = A·s) is the unit of charge (96,485 C are in a mole of singly
charged species = Faraday constant, F ≈ 96,500 C/mol ≈ 105 C/mol)
•
Electricity is the flow of current (A = C/s) and is negative (cathodic) or positive (anodic)
depending on the direction and sign of the current-carrying species (e.g. e–, H+)
•
(Electrode) (electric) potential (V or E; in units of V = J/C) is written as a reduction
This relates to Gibbs free energy as ΔG = –RT ln K = –nFEcell (electrical work per mole), and…
… partial molar Gibbs free energy is the electrochemical potential (𝜇, in units of J/mol)

Chemical potential (𝜇, in units of J/mol)

Galvani/Inner (electric) potential (ϕ, in units of V)
… and in summary, 𝜇 = 𝜇 + 𝑧𝐹ϕ
Also, standard state is a solvent, a solid, and a species at unit activity (~1 M solutes, ~1 bar gases)
Also, E = IR (Ohm’s law) when resistance is constant
•
Galvanic cells produce power (in units of W = A x V = C/s x J/C = J/s) by spontaneous
redox reactions
•
Electrolytic cells require a power input to drive redox reactions; thus, the reactions are
thermodynamically unfavorable
•
Batteries have anodes and cathodes, but these names change depending on if the
battery is being discharged (galvanic) or charged
(electrolytic),
and so
the terminology
(Sources:
B&F, M3LC
course
textbook,
of negative electrode and positive electrode is preferred
and http://goldbook.iupac.org/)
International Union of Pure and Applied
Chemistry (IUPAC)
82
(Accepted) Nomenclature and Terminology that you know, but may have forgotten
•
Coulomb (in units of C = A·s) is the unit of charge (96,485 C are in a mole of singly
charged species = Faraday constant, F ≈ 96,500 C/mol ≈ 105 C/mol)
•
Electricity is the flow of current (A = C/s) and is negative (cathodic) or positive (anodic)
depending on the direction and sign of the current-carrying species (e.g. e–, H+)
•
(Electrode) (electric) potential (V or E; in units of V = J/C) is written as a reduction
This relates to Gibbs free energy as ΔG = –RT ln K = –nFEcell (electrical work per mole), and…
… partial molar Gibbs free energy is the electrochemical potential (𝜇, in units of J/mol)

Chemical potential (𝝁, in units of J/mol)

Galvani/Inner (electric) potential (ϕ, in units of V)
… and in summary, 𝜇 = 𝜇 + 𝑧𝐹ϕ
Also, standard state is a solvent, a solid, and a species at unit activity (~1 M solutes, ~1 bar gases)
Also, E = IR (Ohm’s law) when resistance is constant
Cannot be
measured
independently!
•
Galvanic cells produce power (in units of W = A x V = C/s x J/C = J/s) by spontaneous
redox reactions
•
Electrolytic cells require a power input to drive redox reactions; thus, the reactions are
thermodynamically unfavorable
•
Batteries have anodes and cathodes, but these names change depending on if the
battery is being discharged (galvanic) or charged
(electrolytic),
and so
the terminology
(Sources:
B&F, M3LC
course
textbook,
of negative electrode and positive electrode is preferred
and http://goldbook.iupac.org/)
International Union of Pure and Applied
Chemistry (IUPAC)
83
(Accepted) Nomenclature and Terminology that you know, but may have forgotten
•
Coulomb (in units of C = A·s) is the unit of charge (96,485 C are in a mole of singly
charged species = Faraday constant, F ≈ 96,500 C/mol ≈ 105 C/mol)
•
Electricity is the flow of current (A = C/s) and is negative (cathodic) or positive (anodic)
depending on the direction and sign of the current-carrying species (e.g. e–, H+)
•
(Electrode) (electric) potential (V or E; in units of V = J/C) is written as a reduction
This relates to Gibbs free energy as ΔG = –RT ln K = –nFEcell (electrical work per mole), and…
… partial molar Gibbs free energy is the electrochemical potential (𝜇, in units of J/mol)

Chemical potential (𝜇, in units of J/mol)

Galvani/Inner (electric) potential (ϕ, in units of V)
… and in summary, 𝜇 = 𝜇 + 𝑧𝐹ϕ
Also, standard state is a solvent, a solid, and a species at unit activity (~1 M solutes, ~1 bar gases)
Also, E = IR (Ohm’s law) when resistance is constant
•
Galvanic cells produce power (in units of W = A x V = C/s x J/C = J/s) by spontaneous
redox reactions
•
Electrolytic cells require a power input to drive redox reactions; thus, the reactions are
thermodynamically unfavorable
•
Batteries have anodes and cathodes, but these names change depending on if the
battery is being discharged (galvanic) or charged
(electrolytic),
and so
the terminology
(Sources:
B&F, M3LC
course
textbook,
of negative electrode and positive electrode is preferred
and http://goldbook.iupac.org/)
International Union of Pure and Applied
Chemistry (IUPAC)
84
(Accepted) Nomenclature and Terminology that you know, but may have forgotten
•
Coulomb (in units of C = A·s) is the unit of charge (96,485 C are in a mole of singly
charged species = Faraday constant, F ≈ 96,500 C/mol ≈ 105 C/mol)
•
Electricity is the flow of current (A = C/s) and is negative (cathodic) or positive (anodic)
depending on the direction and sign of the current-carrying species (e.g. e–, H+)
•
(Electrode) (electric) potential (V or E; in units of V = J/C) is written as a reduction
This relates to Gibbs free energy as ΔG = –RT ln K = –nFEcell (electrical work per mole), and…
… partial molar Gibbs free energy is the electrochemical potential (𝜇, in units of J/mol)

Chemical potential (𝜇, in units of J/mol)

Galvani/Inner (electric) potential (ϕ, in units of V)
… and in summary, 𝜇 = 𝜇 + 𝑧𝐹ϕ
Also, standard state is a solvent, a solid, and a species at unit activity (~1 M solutes, ~1 bar gases)
Also, E = IR (Ohm’s law) when resistance is constant
•
Galvanic cells produce power (in units of W = A x V = C/s x J/C = J/s) by spontaneous
redox reactions
•
Electrolytic cells require a power input to drive redox reactions; thus, the reactions are
thermodynamically unfavorable
•
Batteries have anodes and cathodes, but these names change depending on if the
battery is being discharged (galvanic) or charged
(electrolytic),
and so
the terminology
(Sources:
B&F, M3LC
course
textbook,
of negative electrode and positive electrode is preferred
and http://goldbook.iupac.org/)
International Union of Pure and Applied
Chemistry (IUPAC)
85
(Accepted) Nomenclature and Terminology that you know, but may have forgotten
•
Coulomb (in units of C = A·s) is the unit of charge (96,485 C are in a mole of singly
charged species = Faraday constant, F ≈ 96,500 C/mol ≈ 105 C/mol)
•
Electricity is the flow of current (A = C/s) and is negative (cathodic) or positive (anodic)
depending on the direction and sign of the current-carrying species (e.g. e–, H+)
•
(Electrode) (electric) potential (V or E; in units of V = J/C) is written as a reduction
This relates to Gibbs free energy as ΔG = –RT ln K = –nFEcell (electrical work per mole), and…
… partial molar Gibbs free energy is the electrochemical potential (𝜇, in units of J/mol)

Chemical potential (𝜇, in units of J/mol)

Galvani/Inner (electric) potential (ϕ, in units of V)
… and in summary, 𝜇 = 𝜇 + 𝑧𝐹ϕ
Also, standard state is a solvent, a solid, and a species at unit activity (~1 M solutes, ~1 bar gases)
Also, E = IR (Ohm’s law) when resistance is constant
•
Galvanic cells produce power (in units of W = A x V = C/s x J/C = J/s) by spontaneous
redox reactions
opposites
•
Electrolytic cells require a power input to drive redox reactions; thus, the reactions are
thermodynamically unfavorable
•
Batteries have anodes and cathodes, but these names change depending on if the
battery is being discharged (galvanic) or charged
(electrolytic),
and so
the terminology
(Sources:
B&F, M3LC
course
textbook,
of negative electrode and positive electrode is preferred
and http://goldbook.iupac.org/)
International Union of Pure and Applied
Chemistry (IUPAC)
86
(Accepted) Nomenclature and Terminology that you know, but may have forgotten
•
Coulomb (in units of C = A·s) is the unit of charge (96,485 C are in a mole of singly
charged species = Faraday constant, F ≈ 96,500 C/mol ≈ 105 C/mol)
•
Electricity is the flow of current (A = C/s) and is negative (cathodic) or positive (anodic)
depending on the direction and sign of the current-carrying species (e.g. e–, H+)
•
(Electrode) (electric) potential (V or E; in units of V = J/C) is written as a reduction
This relates to Gibbs free energy as ΔG = –RT ln K = –nFEcell (electrical work per mole), and…
… partial molar Gibbs free energy is the electrochemical potential (𝜇, in units of J/mol)

Chemical potential (𝜇, in units of J/mol)

Galvani/Inner (electric) potential (ϕ, in units of V)
… and in summary, 𝜇 = 𝜇 + 𝑧𝐹ϕ
Also, standard state is a solvent, a solid, and a species at unit activity (~1 M solutes, ~1 bar gases)
Also, E = IR (Ohm’s law) when resistance is constant
•
Galvanic cells produce power (in units of W = A x V = C/s x J/C = J/s) by spontaneous
redox reactions
•
Electrolytic cells require a power input to drive redox reactions; thus, the reactions are
thermodynamically unfavorable
•
Batteries have anodes and cathodes, but these names change depending on if the
battery is being discharged (galvanic) or charged (electrolytic), and so the terminology
of negative electrode and positive electrode is preferred
A Very Brief (more rigorous) Review of Thermodynamics
87
FOR YOUR REFERENCE
Electrochemical potential of species i in phase β is an energy (J/mol),
β
𝜇𝑖
=
𝜕𝐺
=
β
𝜕𝑛𝑖
β
𝜇𝑖
+ 𝑧𝑖 𝐹𝜙 β , where
𝑇,𝑝,𝑛𝑗≠𝑖
G (Gibbs free energy (J))
ni (amount of species i (mol))
𝜇𝑖 = 𝜇𝑖 0 + 𝑅𝑇 ln 𝑎𝑖 (chemical potential (J/mol))
zi (valency of species i)
F ≈ 105 (Faraday constant (C/mol)
ϕβ (Galvani/inner electric potential (V))
ai (activity of species i)
β
β
For an uncharged species 𝜇𝑖 = 𝜇𝑖 .
Parsons, Pure & Appl. Chem., 1973, 37, 501
IUPAC Gold (http://goldbook.iupac.org)
Half reactions, at non-unity activity, obey the Nernst equation…
Take ∆𝐺 = ∆𝐺 0 + 𝑅𝑇 ln 𝑄 and use the relation ∆𝐺 = −𝑛𝐹𝐸,
What is Q?
𝑄=
𝑄=
𝑣𝑝
𝑝 𝑎𝑝
𝑣
𝑟 𝑎𝑟 𝑟
𝑣𝑝
𝑝 𝑐𝑝
𝑣
𝑟 𝑐𝑟 𝑟
=
𝑐𝑝
γ
𝑝 𝑝𝑐 0
𝑝
𝑟
𝑐
γ𝑟 𝑟0
𝑣𝑝
𝑣𝑟
, because 𝜇𝑖 = 𝜇𝑖 0 + 𝑅𝑇 ln 𝑎𝑖
𝑐𝑟
, for dilute solutions… which we never have, and where
ap is the activity of product p
ar is the activity of reactant r
vi is the stoichiometric number of i
γi is the activity coefficient of i
ci is the concentration of i
ci0 is the standard state concentration of i
88
Half reactions, at non-unity activity, obey the Nernst equation…
Take ∆𝐺 = ∆𝐺 0 + 𝑅𝑇 ln 𝑄 and use the relation ∆𝐺 = −𝑛𝐹𝐸,
But first… what is Q, again? … the reaction quotient!
𝑄=
𝑄=
𝑣𝑝
𝑝 𝑎𝑝
𝑣
𝑟 𝑎𝑟 𝑟
𝑣𝑝
𝑝 𝑐𝑝
𝑣
𝑟 𝑐𝑟 𝑟
=
𝑐𝑝
γ
𝑝 𝑝𝑐 0
𝑝
𝑟
𝑐
γ𝑟 𝑟0
𝑣𝑝
𝑣𝑟
, because 𝜇𝑖 = 𝜇𝑖 0 + 𝑅𝑇 ln 𝑎𝑖
𝑐𝑟
, for dilute solutions… which we never have, and where
ap is the activity of product p
ar is the activity of reactant r
vi is the stoichiometric number of i
γi is the activity coefficient of i
ci is the concentration of i
ci0 is the standard state concentration of i
89
Half reactions, at non-unity activity, obey the Nernst equation…
Take ∆𝐺 = ∆𝐺 0 + 𝑅𝑇 ln 𝑄 and use the relation ∆𝐺 = −𝑛𝐹𝐸,
But first… what is Q, again? … the reaction quotient!
𝑄=
𝑄=
𝑣𝑝
𝑝 𝑎𝑝
𝑣
𝑟 𝑎𝑟 𝑟
𝑣𝑝
𝑝 𝑐𝑝
𝑣
𝑟 𝑐𝑟 𝑟
=
𝑐𝑝
γ
𝑝 𝑝𝑐 0
𝑝
𝑟
𝑐
γ𝑟 𝑟0
𝑣𝑝
𝑣𝑟
, because fundamentally 𝜇𝑖 = 𝜇𝑖 0 + 𝑅𝑇 ln 𝑎𝑖
𝑐𝑟
, for dilute solutions… which we never have!
ap is the activity of product p
ar is the activity of reactant r
vi is the stoichiometric number of i
γi is the activity coefficient of i
ci is the concentration of i
ci0 is the standard state concentration of i
90
Half reactions, at non-unity activity, obey the Nernst equation…
91
Take ∆𝐺 = ∆𝐺 0 + 𝑅𝑇 ln 𝑄 and use the relation ∆𝐺 = −𝑛𝐹𝐸,
−𝑛𝐹𝐸 =
−𝑛𝐹𝐸 0
+ 𝑅𝑇 ln 𝑄
𝑅𝑇
0
𝐸=𝐸 −
ln 𝑄
𝑛𝐹
Reaction quotient as
product and quotient of
species’ activities
Physicist
Walther Hermann Nernst
(1864–1941)
Nobel Prize (Chemistry, 1920)
from Wiki
Half reactions, at non-unity activity, obey the Nernst equation…
92
Take ∆𝐺 = ∆𝐺 0 + 𝑅𝑇 ln 𝑄 and use the relation ∆𝐺 = −𝑛𝐹𝐸,
−𝑛𝐹𝐸 =
−𝑛𝐹𝐸 0
+ 𝑅𝑇 ln 𝑄
𝑅𝑇
0
𝐸=𝐸 −
ln 𝑄
𝑛𝐹
𝑅𝑇 log 𝑄
0
𝐸=𝐸 −
𝑛𝐹 log 𝑒
𝑅𝑇
0
𝐸=𝐸 −
log 𝑄
0.4343𝑛𝐹
2.3026𝑅𝑇
0
𝐸=𝐸 −
log 𝑄
𝑛𝐹
Reaction quotient as
product and quotient of
species’ activities
Physicist
Walther Hermann Nernst
(1864–1941)
Nobel Prize (Chemistry, 1920)
from Wiki
Half reactions, at non-unity activity, obey the Nernst equation…
93
Take ∆𝐺 = ∆𝐺 0 + 𝑅𝑇 ln 𝑄 and use the relation ∆𝐺 = −𝑛𝐹𝐸,
−𝑛𝐹𝐸 =
−𝑛𝐹𝐸 0
+ 𝑅𝑇 ln 𝑄
𝑅𝑇
0
𝐸=𝐸 −
ln 𝑄
𝑛𝐹
𝑅𝑇 log 𝑄
0
𝐸=𝐸 −
𝑛𝐹 log 𝑒
𝑅𝑇
0
𝐸=𝐸 −
log 𝑄
0.4343𝑛𝐹
2.3026𝑅𝑇
0
𝐸=𝐸 −
log 𝑄
𝑛𝐹
… and at 298.15 K, 𝐸 = 𝐸 0 −
0.05916 V
log 𝑄
𝑛
Reaction quotient as
product and quotient of
species’ activities
Physicist
Walther Hermann Nernst
(1864–1941)
Nobel Prize (Chemistry, 1920)
from Wiki
Half reactions, at non-unity activity, obey the Nernst equation…
94
Take ∆𝐺 = ∆𝐺 0 + 𝑅𝑇 ln 𝑄 and use the relation ∆𝐺 = −𝑛𝐹𝐸,
−𝑛𝐹𝐸 =
−𝑛𝐹𝐸 0
+ 𝑅𝑇 ln 𝑄
𝑅𝑇
0
𝐸=𝐸 −
ln 𝑄
𝑛𝐹
𝑅𝑇 log 𝑄
0
𝐸=𝐸 −
𝑛𝐹 log 𝑒
𝑅𝑇
0
𝐸=𝐸 −
log 𝑄
0.4343𝑛𝐹
2.3026𝑅𝑇
0
𝐸=𝐸 −
log 𝑄
𝑛𝐹
… and at 298.15 K, 𝐸 = 𝐸 0 −
0.05916 V
log 𝑄
𝑛
Reaction quotient as
product and quotient of
species’ activities
Physicist
Walther Hermann Nernst
(1864–1941)
Nobel Prize (Chemistry, 1920)
from Wiki
Remember ~60 mV for log10, but do not forget n and that this is at 25 °C!
95
standard
(SHE)
Rigorously, each needs
to be divided by the
standard-state condition
96
standard
(SHE)
Thus, the potentials for half-cell reactions are actually
full-cell (electric) potential (difference(s)) versus SHE!
97
standard
(SHE)
* Normal hydrogen electrode (NHE) is an empirical SHE ([H+] = 1; not standard state)
* Standard hydrogen electrode (SHE) is a hypothetical, perfect NHE (a = 1; not empirical)
* Reversible hydrogen electrode (RHE) is the SHE but the same regardless of pH
* And generally, formal potentials (E0’) take into consideration non-idealities and changes
in ionic strengths so that the reaction quotient only has concentrations, and not activities
Ramette, J. Chem. Educ., 1987, 64, 885
98
standard
(SHE)
* Normal hydrogen electrode (NHE) is an empirical SHE ([H+] = 1; not standard state)
* Standard hydrogen electrode (SHE) is a hypothetical, perfect NHE (a = 1; not empirical)
* Reversible hydrogen electrode (RHE) is the SHE but the same regardless of pH
* And generally, formal potentials (E0’) take into consideration non-idealities and changes
in ionic strengths so that the reaction quotient only has concentrations, and not activities
Potential
Standard Potential
Ramette, J. Chem. Educ., 1987, 64, 885
99
standard
(SHE)
Given E0 = 0, what is
this experimental redox
potential versus?
Potential
Standard Potential
100
standard
(SHE)
Given E0 = 0, what is
this experimental redox
potential versus?
vs. SHE
Potential
Standard Potential
101
EXAMPLE: Write a balanced chemical equation and
calculate the standard cell potential for the galvanic cell:
Zn(s) | Zn2+ (1 M) || MnO4– (1 M), Mn2+ (1 M), H+ (1 M) | Pt(s)
Look up half-reactions and standard reduction potentials in an
Electrochemical Series table (CRC, B&F Appendix C, WWW):
Anode:
Zn
Zn2+ + 2e–
E0anode = -0.76 V
Note: Although strictly correct, do not use “–E0” as the “E0 for oxidation”
Cathode: MnO4– + 8H+ + 5e–
Mn2+ + 4H2O E0cathode = +1.51 V
Note: Be careful to choose the correct half-reaction with MnO4–
CRC Handbook of Chemistry and Physics, 92nd Edition
102
* All values versus SHE
MnO4– + 8H+ + 5e–
Mn2+ + 4H2O
Zn
Zn2+ + 2e–
http://folk.ntnu.no/andersty/2.%20Klasse/KJ1042%20Termodynamikk%20med%20lab/Lab/
Oppgave%205%20-%20Standard%20reduksjonspotensial/Rapportfiler/E0.pdf
103
EXAMPLE: Write a balanced chemical equation and
calculate the standard cell potential for the galvanic cell:
Zn(s) | Zn2+ (1 M) || MnO4– (1 M), Mn2+ (1 M), H+ (1 M) | Pt(s)
Look up half-reactions and standard reduction potentials in an
Electrochemical Series table (CRC, B&F Appendix C, WWW):
Anode:
Zn
Zn2+ + 2e–
E0anode = -0.76 V
Note: Although strictly correct, do not use “–E0” as the “E0 for oxidation”
Cathode: MnO4– + 8H+ + 5e–
Mn2+ + 4H2O E0cathode = +1.51 V
Note: Be careful to choose the correct half-reaction with MnO4–
To get the balanced overall reaction… ?
104
EXAMPLE: Write a balanced chemical equation and
calculate the standard cell potential for the galvanic cell:
Zn(s) | Zn2+ (1 M) || MnO4– (1 M), Mn2+ (1 M), H+ (1 M) | Pt(s)
Look up half-reactions and standard reduction potentials in an
Electrochemical Series table (CRC, B&F Appendix C, WWW):
Anode:
Zn
Zn2+ + 2e–
E0anode = -0.76 V
Note: Although strictly correct, do not use “–E0” as the “E0 for oxidation”
Cathode: MnO4– + 8H+ + 5e–
Mn2+ + 4H2O E0cathode = +1.51 V
Note: Be careful to choose the correct half-reaction with MnO4–
To get the balanced overall reaction, multiply the anode
reaction by 5 and add to 2 times the cathode reaction, giving:
2MnO4– (aq) + 16H+ (aq) + 5Zn (s)
5Zn2+ (aq) + 2Mn2+ (aq) + 8H2O (l)
E0cell = E0cathode – E0anode = 1.51 V + 0.76 V = +2.27 V
Q: What is electrochemistry? … one more … (remember this?)
105
CRC Handbook of Chemistry and Physics, 92nd Edition
* All values versus SHE
* How negative can Eo be?
http://folk.ntnu.no/andersty/2.%20Klasse/KJ1042%20Termodynamikk%20med%20lab/Lab/
Oppgave%205%20-%20Standard%20reduksjonspotensial/Rapportfiler/E0.pdf
Q: What is electrochemistry? … one more …
A: Any process involving the motion/transport of charge – carried
by entities other than unsolvated electrons and holes – through
phase(s), or the transfer of charge across interface(s).
106
Example: solvated electrons
Prof. Robert Hamers (Univ. of Wisconsin)
http://hamers.chem.wisc.edu/people
diamond
Zhu, …, Hamers, Nature Materials, 2013, 12, 836
First description of conductivity using solvated electrons
107
During the first part of the twentieth century, E. C. Franklin and C. A. Kraus probably did more to elucidate the chemistry of
liquid ammonia solutions than everybody else combined… It is perhaps little known that their work was prompted by the
research and insight of H. P. Cady, carried out while he was an undergraduate! Whilst working on cobalt ammine
complexes, Cady proposed that ammonia in these (and other “double salts”) must function in a manner akin to water in salts
with water of crystallization. He suggested further that liquid ammonia would probably be found to resemble water in its
physical and chemical properties—thus adding a second to our list of ionizing solvents. Cady’s undergraduate work, carried
out without supervision, published in 1897, was perhaps the first physical chemistry study of liquid ammonia solutions.
Zurek, Edwards, & Hoffman, Angew. Chem. Int. Ed., 2009, 48, 8198
Cady, J. Phys. Chem., 1897, 1, 707
Q: What is electrochemistry? … one more …
A: Any process involving the motion/transport of charge – carried
by entities other than unsolvated electrons and holes – through
phase(s), or the transfer of charge across interface(s).
108
Example: solvated electrons
Prof. Robert Hamers (Univ. of Wisconsin)
http://hamers.chem.wisc.edu/people
What is this?
diamond
Zhu, …, Hamers, Nature Materials, 2013, 12, 836
Absolute potentials can be measured / approximated…
very carefully…
109
Prof. Sergio Trasatti
(Università de Milano, Italy)
Trasatti, Pure & Appl. Chem., 1986, 58, 955
Absolute potentials can be measured / approximated…
very carefully…
FOR YOUR REFERENCE
110
Trasatti, Pure & Appl. Chem., 1986, 58, 955
Absolute potentials can be measured / approximated…
very carefully…
FOR YOUR REFERENCE
111
Born–Haber cycle:
𝐻𝑔
𝜇𝑒
e(Hg)
Trasatti, Electroanal. Chem. Interfac. Electrochem., 1974, 52, 313
Trasatti, Pure & Appl. Chem., 1986, 58, 955
Absolute potentials can be measured / approximated…
very carefully…
FOR YOUR REFERENCE
112
Born–Haber cycle:
𝐻𝑔
𝜇𝑒
e(Hg)
… but we need an
Trasatti, Electroanal. Chem. Interfac. Electrochem., 1974, 52, 313
electrode!
Trasatti, Pure & Appl. Chem., 1986, 58, 955
Absolute potentials can be measured / approximated…
very carefully…
FOR YOUR REFERENCE
113
Born–Haber cycle:
𝐻𝑔
𝜇𝑒
e(Hg)
… but we need an
Trasatti, Electroanal. Chem. Interfac. Electrochem., 1974, 52, 313
electrode!
Trasatti, Pure & Appl. Chem., 1986, 58, 955
Absolute potentials can be measured / approximated…
very carefully…
FOR YOUR REFERENCE
114
Born–Haber cycle:
𝐻𝑔
𝜇𝑒
e(Hg)
… but we need an
Trasatti, Electroanal. Chem. Interfac. Electrochem., 1974, 52, 313
electrode!
Trasatti, Pure & Appl. Chem., 1986, 58, 955
Absolute potentials can be measured / approximated…
very carefully…
FOR YOUR REFERENCE
115
Born–Haber cycle:
𝐻𝑔
𝜇𝑒
e(Hg)
… but we need an
Trasatti, Electroanal. Chem. Interfac. Electrochem., 1974, 52, 313
electrode!
Trasatti, Pure & Appl. Chem., 1986, 58, 955
Absolute potentials can be measured / approximated…
very carefully…
FOR YOUR REFERENCE
116
Hg
Pt
Born–Haber cycle:
𝐻𝑔
𝜇𝑒
e(Hg)
… but we need an
Farrell and McTigue, Electroanal. Chem. Interfac. Electrochem., 1982, 139, 37
Trasatti, Electroanal. Chem. Interfac. Electrochem., 1974, 52, 313
electrode!
Trasatti, Pure & Appl. Chem., 1986, 58, 955
Two diagrams of empirical standard potentials…
117
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
Chemist
from Wiki
Wendell Mitchell Latimer
(1893–1955)
http://academictree.org/chemistry/peopleinfo.php?pid=24644
Latimer, The oxidation states of the elements and their potentials in aqueous solution, 1938
Two diagrams of empirical standard potentials…
118
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
7+
6+
4+
3+
2+
0
Chemist
from Wiki
Wendell Mitchell Latimer
(1893–1955)
http://academictree.org/chemistry/peopleinfo.php?pid=24644
Latimer, The oxidation states of the elements and their potentials in aqueous solution, 1938
Two diagrams of empirical standard potentials…
119
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
7+
6+
4+
3+
2+
0
Chemist
from Wiki
Disproportionation – spontaneous and simultaneous
reduction and oxidation of a molecule
Wendell Mitchell Latimer
(1893–1955)
http://academictree.org/chemistry/peopleinfo.php?pid=24644
Latimer, The oxidation states of the elements and their potentials in aqueous solution, 1938
Two diagrams of empirical standard potentials…
120
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
7+
6+
4+
3+
2+
0
from Wiki
Disproportionation – spontaneous and simultaneous
reduction and oxidation of a molecule
(1) Does Mn2+ disproportionate?
(2) What is the standard reduction potential of MnO4– to MnO2?
Two diagrams of empirical standard potentials…
121
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
7+
6+
4+
3+
2+
0
from Wiki
Disproportionation – spontaneous and simultaneous
reduction and oxidation of a molecule
(1) Does Mn2+ disproportionate?
(2) What is the standard reduction potential of MnO4– to MnO2?
Reduction: Mn2+
Mno
Eo = +1.18 V
Oxidation: Mn2+
Mn3+
Eo = +1.51 V
Two diagrams of empirical standard potentials…
122
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
7+
6+
4+
3+
2+
0
from Wiki
Disproportionation – spontaneous and simultaneous
reduction and oxidation of a molecule
(1) Does Mn2+ disproportionate?
NO. Eo = Ered – Eox = 1.18 – 1.51 = –0.33 V
(2) What is the standard reduction potential of MnO4– to MnO2?
Reduction: Mn2+
Mno
Eo = +1.18 V
Oxidation: Mn2+
Mn3+
Eo = +1.51 V
Two diagrams of empirical standard potentials…
123
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
7+
6+
4+
3+
2+
0
from Wiki
Disproportionation – spontaneous and simultaneous
reduction and oxidation of a molecule
(1) Does Mn2+ disproportionate?
NO. Eo = Ered – Eox = 1.18 – 1.51 = –0.33 V
(2) What is the standard reduction potential of MnO4– to MnO2?
ΔGo = -nFEo = -3FEo
Two diagrams of empirical standard potentials…
124
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
7+
6+
4+
3+
2+
0
from Wiki
Disproportionation – spontaneous and simultaneous
reduction and oxidation of a molecule
(1) Does Mn2+ disproportionate?
NO. Eo = Ered – Eox = 1.18 – 1.51 = –0.33 V
(2) What is the standard reduction potential of MnO4– to MnO2?
ΔGo = -nFEo = -3FEo
ΔGo = -nFEo1 + -nFEo2 = -F((1 x 0.56 V) + (2 x 2.26 V)) = -F(5.08 V)
Two diagrams of empirical standard potentials…
125
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
7+
6+
4+
3+
2+
0
from Wiki
Disproportionation – spontaneous and simultaneous
reduction and oxidation of a molecule
(1) Does Mn2+ disproportionate?
NO. Eo = Ered – Eox = 1.18 – 1.51 = –0.33 V
(2) What is the standard reduction potential of MnO4– to MnO2?
ΔGo = -nFEo = -3FEo
ΔGo = -nFEo1 + -nFEo2 = -F((1 x 0.56 V) + (2 x 2.26 V)) = -F(5.08 V)
Set them equal to each other, and thus, 3Eo = 5.08 and Eo = 1.69 V
Two diagrams of empirical standard potentials…
126
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
7+
6+
4+
3+
2+
0
from Wiki
Recall from before…
Two diagrams of empirical standard potentials…
127
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
???
7+
6+
4+
3+
2+
0
from Wiki
Recall from before…
???
Two diagrams of empirical standard potentials…
128
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
???
7+
6+
4+
3+
2+
0
from Wiki
Recall from before…
… anyway, why are
these bottom E0
values not on the
Latimer diagram?
???
Two diagrams of empirical standard potentials…
129
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
???
7+
6+
4+
3+
2+
0
from Wiki
Recall from before…
… anyway, why are
these bottom E0
values not on the
Latimer diagram?
???
… because they
are at basic/alkaline
standard state with
~1 M OH–!
Two diagrams of empirical standard potentials…
130
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
???
7+
6+
4+
3+
2+
0
from Wiki
Recall from before…
What would this E0
value be when at
acidic standard
state?
???
Two diagrams of empirical standard potentials…
131
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
FOR YOUR REFERENCE
Reduction
Oxidation
1,69 V
???
7+
6+
4+
3+
2+
0
from Wiki
???
𝐸=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
𝑛
What would this E0
value be when at
acidic standard
state?
𝑀𝑛𝑂2
𝑀𝑛𝑂4
1
𝐻2 𝑂
2− 1
𝐻+
2
4
=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
2
1
1
11
10−14
4
0
= 𝐸𝑎𝑐𝑖𝑑
− 0.02958 V 56
Two diagrams of empirical standard potentials…
132
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
FOR YOUR REFERENCE
Reduction
Oxidation
1,69 V
???
7+
6+
4+
3+
2+
0
from Wiki
???
𝐸=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
𝑛
What would this E0
value be when at
acidic standard
state?
𝑀𝑛𝑂2
𝑀𝑛𝑂4
1
𝐻2 𝑂
2− 1
𝐻+
2
4
=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
2
1
1
11
10−14
4
0
= 𝐸𝑎𝑐𝑖𝑑
− 0.02958 V 56
Two diagrams of empirical standard potentials…
133
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
FOR YOUR REFERENCE
Reduction
Oxidation
1,69 V
???
7+
6+
4+
3+
2+
0
from Wiki
???
𝐸=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
𝑛
What would this E0
value be when at
acidic standard
state?
𝑀𝑛𝑂2
𝑀𝑛𝑂4
1
𝐻2 𝑂
2− 1
𝐻+
2
4
=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
2
1
1
11
10−14
4
0
= 𝐸𝑎𝑐𝑖𝑑
− 0.02958 V 56
Two diagrams of empirical standard potentials…
134
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
FOR YOUR REFERENCE
Reduction
Oxidation
1,69 V
???
7+
6+
4+
3+
2+
0
from Wiki
???
𝐸=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
𝑛
𝑀𝑛𝑂2
𝑀𝑛𝑂4
1
𝐻2 𝑂
2− 1
𝐻+
2
4
=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
2
1
1
11
10−14
4
0
= 𝐸𝑎𝑐𝑖𝑑
− 0.02958 V 56
0
𝐸 = 𝐸𝑎𝑐𝑖𝑑
− 1.65648 V = 0.60 V
What would this E0
value be when at
acidic standard
state?
Two diagrams of empirical standard potentials…
135
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
FOR YOUR REFERENCE
Reduction
Oxidation
1,69 V
???
7+
6+
4+
3+
2+
0
from Wiki
???
𝐸=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
𝑛
𝑀𝑛𝑂2
𝑀𝑛𝑂4
1
𝐻2 𝑂
2− 1
𝐻+
2
4
=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
2
1
1
11
10−14
4
0
= 𝐸𝑎𝑐𝑖𝑑
− 0.02958 V 56
0
𝐸 = 𝐸𝑎𝑐𝑖𝑑
− 1.65648 V = 0.60 V
What would this E0
value be when at
acidic standard
state?
0
𝐸𝑆𝐻𝐸
= 2,25648 V
SWEET!
Two diagrams of empirical standard potentials…
136
A Latimer diagram is a summary of the E0 values for an element; it is useful for visualizing the
complete redox series for an element and for determining when disproportionation will occur.
Reduction
Oxidation
1,69 V
???
7+
6+
4+
3+
2+
0
from Wiki
???
𝐸=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
𝑛
𝑀𝑛𝑂2
𝑀𝑛𝑂4
1
𝐻2 𝑂
2− 1
𝐻+
2
4
=
0
𝐸𝑎𝑐𝑖𝑑
0.05916 V
−
log
2
1
1
11
10−14
4
0
= 𝐸𝑎𝑐𝑖𝑑
− 0.02958 V 56
0
𝐸 = 𝐸𝑎𝑐𝑖𝑑
− 1.65648 V = 0.60 V
What would this E0
value be when at
acidic standard
state?
0
𝐸𝑆𝐻𝐸
= 2,25648 V
SWEET!
… but then why did the
CRC not list this? …
… Second one (not truly standard potentials)…
137
A Pourbaix diagram is a map of the predominant equilibrium species of an aqueous
electrochemical system; it is useful for identifying which materials/species are present/stable
Chemist
Marcel Pourbaix
(1904–1998)
http://corrosion-doctors.org/Biographies/PourbaixBio.htm
from Wiki
Pourbaix, Atlas of electrochemical equilibria in aqueous solutions, 1974
… but then why did the
CRC not list this? …
… Second one (not truly standard potentials)…
138
A Pourbaix diagram is a map of the predominant equilibrium species of an aqueous
electrochemical system; it is useful for identifying which materials/species are present/stable
Chemist
Marcel Pourbaix
(1904–1998)
http://corrosion-doctors.org/Biographies/PourbaixBio.htm
from Wiki
Oxidation
1,69 V
???
… because in acid, the reaction does
not occur!
… but then why did the
from Wiki
CRC not list this? …
Pourbaix, Atlas of electrochemical equilibria in aqueous solutions, 1974
7+
6+
4+
3+
2+
0
… Second one (not truly standard potentials)…
139
A Pourbaix diagram is a map of the predominant equilibrium species of an aqueous
electrochemical system; it is useful for identifying which materials/species are present/stable
Chemist
E(O2,H+/H2O)
Marcel Pourbaix
(1904–1998)
http://corrosion-doctors.org/Biographies/PourbaixBio.htm
Why don’t I like this? …
Even though EVERYONE
does it this way
Anyway, … standard state is here
Pourbaix, Atlas of electrochemical equilibria in aqueous solutions, 1974
RHE
from Wiki