CHEM1612 - Pharmacy
Week 8: Redox Reactions
Dr. Siegbert Schmid
School of Chemistry, Rm 223
Phone: 9351 4196
E-mail: siegbert.schmid@sydney.edu.au
Unless otherwise stated, all images in this file have been reproduced from:
Blackman, Bottle, Schmid, Mocerino and Wille,
Chemistry, John Wiley & Sons Australia, Ltd. 2008
ISBN: 9 78047081 0866
Electrochemistry
n
Blackman, Bottle, Schmid, Mocerino & Wille:
Chapter 12, Sections 4.8 and 4.9
Key chemical concepts:
n  Redox and half reactions
n  Cell potential
n  Voltaic and electrolytic cells
n  Concentration cells
Key Calculations:
n  Calculating cell potential
n  Calculating amount of product for given current
n  Using the Nernst equation for concentration cells
Lecture 25-3
The Electron
n
n
n
n
n
Symbol: emass = 9.109 · 10–31 kg
Possesses unit negative charge = -1.602 · 10–19 C ( 1 C = 1 J V-1)
In a bond it is attracted to more electronegative atoms
Given Coulomb’s law:
q
q
It is attracted towards positive charges.
It is repelled by negative charges.
Lecture 25-4
Reductions and Oxidations
n
n
Mn7+ à Mn2+
Mn(VII) à Mn (II)
Mn has gained electrons in this reaction.
We say that Mn has been reduced.
n
n
Cl- à Cl2
Cl(–I) à Cl (0)
Electrons have left the chloride.
We say that Cl- has been oxidised.
Lecture 25-5
Redox Reactions
n
A REDOX reaction (or REDuction – OXidation reaction) involves
the transfer of electrons
from one element (the reducing agent)
to another (the oxidising agent).
n
Involves a change in oxidation state (number) of the involved
species.
e.g.
n
n
Zn(s) + Cu2+(aq) à Zn2+(aq) + Cu(s)
Two electrons have moved from Zn to Cu2+
Electrochemical processes convert chemical energy to
electrical energy via redox reactions!
Lecture 25-6
Summary of terms for Redox reactions
e-
X
transfer
Y
of
electrons
X loses electron(s)
Y gains electron(s)
X increases its
oxidation number
Y decreases its
oxidation number
X is the reducing agent
Y is the oxidizing agent
X is oxidized
Y is reduced
Lecture 25-7
Not-redox reactions
n
These are not redox reactions because the O.N. does not change:
H2O + CH3COOCH3 à CH3COOH + CH3OH
Ag+ (aq) + Cl– (aq) à AgCl (s)
What affects the redox behaviour of an atom?
Balance between Electron Affinity and Ionisation Energy!!
Lecture 25-8
Ionisation Energy
n
Ionisation Energy = Energy required to remove an electron from an
atom or ion
M(g) à M+(g)(kJ mol–1) + e-
n
n
oxidation
It is low for elements that tend to form cations
Na
Mg
Al
Si
P
S
Cl
Ar
496
738
578
787
1012
1000
1250
1520
Atoms with high I.E. (except noble gases) often form anions
Lecture 25-9
Electron Affinities
n
Electron Affinity = Energy released when
an electron associates with an atom/ion
X(g) + e- à X–(g) (kJ mol–1)
B
–15
n
C
–121
N
31
O
–142
reduction
F
–333
The more negative the electron affinity the more energy is released
upon addition of an electron, i.e. e- addition is more favourable.
Lecture 25-10
Redox Reaction: Example
n
In a redox reaction, there is a change in the O.N. of the species
involved:
e.g. smelting of iron
oxidising
agent
reducing
agent
Fe2O3 (s) + 3 CO (g) → 2 Fe (l) + 3 CO2(g)
+3 -2
+2 -2
Reduced
Total = 2 ·(-3)
0
+4 -2
Oxidised
Total = 3 · (+2)
Blast Furnace
Lecture 25-11
Half-reactions
Example of redox reaction:
n
Zn(s) + Cu2+(aq) à Zn2+(aq) + Cu(s)
Any redox reaction can be broken down into two separate halfreactions, showing the process of reduction and oxidation in isolation.
Zn (s) à Zn2+ (aq) + 2e– (aq)
OXIDATION – loss of
electrons
Half reaction
Cu2+ (aq) + 2e– (aq) à Cu (s)
REDUCTION – gain of
electrons
Lecture 25-12
Activity series of metals
Consider the same reaction for 5 metals
(oxidation of the metal by hydrogen ions):
Observations:
Cu(s) + 2H+(aq) à Cu2+(aq) + H2(g)
no reaction
Sn(s) + 2H+(aq) à Sn2+(aq) + H2(g)
no reaction
Fe(s) + 2H+(aq) à Fe2+(aq) + H2(g)
yellow colour slowly emerges
Zn(s) + 2H+(aq) à Zn2+(aq) + H2(g)
gas bubbling
Mg(s) + 2H+(aq) à Mg2+(aq) + H2(g)
vigorous reaction, Mg disappears
Activity: Mg > Zn > Fe > Sn > Cu as reducing agents
Lecture 25-13
Activity series of metals
We just saw activity of metals, compared to H+:
Mg > Zn > Fe > Sn > Cu
(1)
Zn(s) + 2H+(aq) à Zn2+(aq) + H2(g)
gas bubbling
(2)
Cu(s) + 2H+(aq) à Cu2+(aq) + H2(g)
no reaction
(1) - (2)
Zn(s) + Cu2+(aq) D Zn2+(aq) + Cu(s)
We immerse in a blue solution of CuSO4 a strip of metallic Zn.
Observations:
What
will happen?
Cu precipitates on Zn - the
solution
will slowly change color
Reason: Zn is a stronger reducing agent than Cu, so reaction will
proceed to the right.
Lecture 25-14
The molecular interpretation
Cu2+ solution
Zinc strip
Layer of red-brown
copper
Lecture 25-15
Half-reactions
In this redox reaction occurring in a beaker the electrons are directly
transferred from one element to another.
Zn à Zn2+ + 2eCu2+ + 2e- à Cu
Zn(s) + Cu2+(aq) à Zn2+(aq) + Cu(s)
Note: phase
labels do not
appear in halfreactions
This means that we cannot access the electrons produced in the
redox process.
n
Lecture 25-16
Galvanic cells / Voltaic Cells
n
If we can separate the half-reactions then we can harness the
electrons. The redox pair can be kept separate, with electrontransfer taking place through an external circuit.
n
The separation of half reactions is the basis of an electrochemical
cell, a device capable of transforming electric current into energy
and viceversa.
n
A cell that uses the energy released by a spontaneous redox
reaction to produce energy is known as galvanic or voltaic cell:
(a battery: the spontaneous reaction generates electricity).
Images
Luigi
Galvani
(1737-1798)
Count
Alessandro
Volta (1745-1827)
downloaded
from
Wikipedia
Lecture 25-17
Galvanic (or Voltaic) Cell
Ag+(aq) + ½Cu(s) à Ag(s) + ½Cu2+(aq)
Lecture 25-18
Electromotive Force
n
A voltaic cell involves two half cells - one containing an oxidising agent
and the other a reducing agent.
n
These cells are connected with a wire, to allow electron flow.
n
The measured voltage across the cell, the cell potential (Ecell) is called
the ELECTROMOTIVE FORCE (emf) of the cell.
Positive emf = Spontaneous Reaction
(does work on the universe)
Lecture 25-19
Voltaic Cell Notation / Line Notation
n
Rather than writing out the full chemical equation, we can use a
short-hand notation:
Eg.
Zn(s) + Cu2+(aq) à Zn2+(aq) + Cu(s)
Salt bridge
Anode half-cell
Cathode half-cell
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Phase boundary
Phase boundary
Electrons flow this way
Lecture 25-20
Challenge Quiz
n
Which of these compounds is likely to be a reducing agent (red) and
which an oxidising agent (ox)?
q
q
q
q
q
q
q
q
q
q
q
Mn2+
Fe3+
Zn
Na+
Cr2O72Cu2+
K
H+
Br2
MnO4Sn2+
ox and red, has O.N. 0,+3,+4,+7
ox;+3 is its higher O.N.
red, has ON +2
ox;+1 is its higher O.N.
ox;+6 is its higher O.N.
ox;+2 is its higher O.N.
red, has O.N. +1
ox;+1 is its higher O.N.
ox; 0 is its higher O.N.
ox; +7 is its higher O.N.
ox and red, has O.N. +4
Lecture 25-21
Electromotive Force
n
n
The measured voltage across the cell is called cell potential (Ecell).
This driving force for the reaction is also called ELECTROMOTIVE
FORCE (emf) of the cell.
Figure from Silberberg,
“Chemistry”,
McGraw Hill, 2006.
~1.1 V
Zn(s) + Cu2+(aq) à Zn2+(aq) + Cu(s)
n
n
n
Every galvanic cell has a different cell potential.
How can we measure the cell potential relative to each species?
We could tabulate the cell potential of each element against a
standard reference electrode.
Lecture 25-22
Summary of Electrochemistry Concepts
q
q
q
q
q
q
q
q
q
q
Redox reactions
Standard reduction potential, E0
Reference electrodes
Galvanic cells, cell notation, and electromotive force Ecell
Electrolytic cells and Faraday’s Law
Nernst Equation and concentration cells
Examples of biological concentration cells
Relationship between E0, ΔG, Q, and K
Corrosion
Batteries
Lecture - 23